





Glass Q J3 A-S’ 
Book .H 7 5Z, 
Gojyriglif}! ? /31 

COPYRIGHT DEPOSIT 


V> 









LABORATORY EXERCISES 

AND 

PROBLEMS IN GENERAL CHEMISTRY 


BY 

B SMITH HOPKINS 

PROFESSOR OF CHEMISTRY, UNIVERSITY OF ILLINOIS 
AND 

HARVEY aV"' NEVILLE 

t 

ASSOCIATE PROFESSOR OF CHEMISTRY, LEHIGH UNIVERSITY 



Revised Edition 


D. C. HEATH AND COMPANY 

BOSTON NEW YORK CHICAGO 

ATLANTA SAN FRANCISCO DALLAS 

LONDON 

C 



^ A 

:*\\ 

1 A 5' ' 

\ ' 

ACKNOWLEDGMENT 

In the development of these exercises the authors are indebted 
to the members of the Division of Inorganic Chemistry of the 
University of Illinois for many helpful suggestions. In general, 
the experiments are the result of the teaching experience in this 
laboratory. 



Copyright, 1931 
By D. C. Heath and Company 

3 1 1 


©CIA 44412 

PRINTED IN U. S. A. 


OCT 24 1931 


CONTENTS 

PART I-THE NON-METALS 

EXERCISE PAGE 

General Laboratory Rules. 1 

Suggestions for Writing Laboratory Notes. 2 

Procedure for the First Laboratory Period. 3 

1. The Use of Glass Tubing. 4 

2. Measurement — The Metric System. 6 

Laboratory Technique. 7 

3. Properties of Matter — Chemical and Physical Change 10 

4. Elements, Compounds — Separation from Mixtures. . . 11 

5. Study of a Mixture. 12 

6. Chemical Change. 13 

7. Sources of Oxygen. 14 

8. Catalysis. 15 

9. Preparation of Oxygen; Combustion — Formation of 

Acids and Bases. 15 

10. Accurate Weighing. 17 

11. Weight of a Liter of Oxygen — Law of Constant Compo¬ 

sition . 19 

12. Methods of Preparing Hydrogen. 21 

13. Preparation and Properties of Hydrogen. 22 

14. Hydrogen as a Reducing Agent. 23 

15. Weight of Metal to Displace a Gram Atom of Hydrogen 24 

16. Water — Boiling and Freezing Points — Purification. . 26 

17. Some Properties of Water—Solutions. 27 

18. Hydrogen Peroxide. 28 

19. The Law of Multiple Proportions. 30 

20. Solubility — Supersaturated Solution. 31 

21. Hydrates. 32 

22. Formula of a Hydrate . 33 

23. Colloidal Solutions. 34 

24. Gels. 36 

25. Double Decomposition or Metathesis. 37 

26. Preparation and Properties of Chlorine. 38 

27. Hydrogen Chloride — Hydrochloric Acid. 39 

28. Potassium Hypochlorite. 41 

29. Potassium Chlorate. 42 

30. Potassium Perchlorate . 42 

31. Bleaching Powder. 43 

32. Ionization. 44 

33. Ionization and Chemical Action. 45 

iii 


































iv CONTENTS 

EXERCISE PAGE 

34. Electrolysis of a Salt Solution. 46 

35. Hydrolysis . 46 

36. Bromine. 47 

37. Iodine. 48 

38. Hydrogen Fluoride (Hydrofluoric Acid). 49 

39. Preparation of Pure Sodium Chloride from Rock Salt 50 

40. Atomic Structure. 51 

41. The Per Cent of Oxygen in the Air. 52 

42. Ammonia — Preparation; Properties. 53 

43. Nitric Acid — Preparation and Properties. 54 

44. Reduction Products of Nitric Acid. 56 

45. Nitrous Oxide. 57 

46. Nitrites — Nitrous Anhydride. 58 

47. Sulfur — Allotropic Forms. 58 

48. Sulfur — Combination with Metals; Hydrogen Sulfide 59 

49. Hydrogen Sulfide — Properties and Uses. 60 

50. Sulfur Dioxide — Preparation; Properties. 61 

51. Sulfuric Acid. 63 

52. Sodium Thiosulfate. 64 

53. Normal Solutions — Equivalent Weights. 65 

54. Normal Solutions — Acidimetry. 66 

55. Phosphorus. 67 

56. Arsenic and Antimony. 69 

57. Bismuth. 70 

58. Silicon . 71 

59. Boron. 72 

60. Carbon. 73 

61. Identification of a Salt. 74 

62. A Review of Oxidation-Reduction. 75 

PART II —THE METALS 

63. The Physical Properties of the Metals. 77 

64. Properties of Alloys. 78 

65. The Chemical Properties of the Metals. 82 

66. Properties of Sodium. 84 

67. Preparation and Properties of Sodium Hydroxide .... 86 

68. Hydroxides of the Metals . 88 

69. Sodium Bicarbonate and Sodium Carbonate. 89 

70. Preparation of Potassium Nitrate. 91 

71. Uses of Some Potassium Compounds. 93 

72. Ammonium Compounds. 94 

73. Tests for Potassium, Sodium, Ammonium Salts. 96 

74. Soap. 97 

75. Emulsions. 99 

76. Cuprous Compounds.101 

77. Preparation of Some Cupric Compounds.102 

78. Metallic Couples.104 











































CONTENTS V 

EXERCISE PAGE 

79. Electrochemical Action.105 

80. Silver.107 

81. Reactions of Calcium, Strontium, and Barium.109 

82. Hard Water .Ill 

83. Magnesium.112 

84. Zinc and Cadmium. 114 

85. Mercury.115 

86 . The Preparation of Mercuric Thiocyanate.117 

87. Aluminum.119 

88 . Qualitative Analysis of Baking Powder.121 

89. Tin.122 

90. Lead .124 

91. Chromium.126 

92. Preparation of Chrome Alum.127 

93. Manganese .129 

94. Iron — Ferric and Ferrous Compounds.130 

95. Corrosion of Iron.131 

96. Preparation of Double Salts.132 

97. Cobalt. 133 

98. Nickel.134 

99. Unknown No. 5 — “Group Unknown”.135 

100. Unknown No. 6 — “General Unknown”.136 

PART III —CHEMICAL ARITHMETIC 

Types of Problems.141 

Appendix. 1^1 


























LIST OF SUPPLIES 

The following is an estimate of the quantities of material 
and apparatus advisable to purchase for a class of ten students 
who perform all the experiments in the manual. The amounts 
listed are generous to allow for waste of chemicals and reason¬ 
able breakage of apparatus. 

In certain experiments requiring the use of the analytical 
balance or other limited equipment it will be found advisable 
not to have all the students perform such experiments in the 
same period. 


GENERAL APPARATUS 

10 asbestos mats, 10 cm. X 10 cm. 

1 balance, analytical, with weights and rider for weighing 
from 0.1 mg. to 100 g. 

2 balances, platform, with weights for weighing from 0.1 g. 
to 1000 g. 

1 barometer 
10 brushes, test-tube 
10 burners, Bunsen 
10 clamps, iron, small 
10 clamps, iron, large 
1 gross corks, assorted sizes 
1 set cork borers, 6 in set 
10 crucibles, iron, diam. 4 cm. 

24 crucibles, porcelain, #00 

10 cylinders, porcelain, unglazed, diam. 3 cm. 

24 dishes, porcelain, evaporating, 8 cm. 

6 electrodes, carbon, diam. 5 mm. 

10 files, triangular, 15 cm. 

10 pr. forceps, iron, 10 cm. 

15 squares gauze, iron wire, asbestos center, 12 cm. X 12 cm. 
10 holders, test-tube 
5 magnets, horse-shoe, small 
5 magnifiers or lenses 


VI 


LIST OF SUPPLIES 


vii 

10 meter sticks, graduated in inches on reverse side 
10 mortars, porcelain, with pestle, 9 cm. 

10 pks. paper, filter, qualitative, 13 cm. 

2 pks. paper, filter, qualitative, 20 cm. 

12 vials each, paper, litmus, red and blue 
12 pinch cocks, Mohr's, medium 
1 m. platinum wire, #25 
10 racks, test-tube 
10 sponges 

10 spoons, iron, deflagrating, diam. of bowl 1 cm. 

10 spoons, porcelain, spatula end, 15 cm. 

10 stands, filter 

10 stands, iron, with 2 rings 

1 kg. stoppers, rubber, assorted sizes #0-6, one- and two-hole 
10 thermometers, chemical, 0°-250° C. 

10 triangles, pipestem, to support crucible #00 
10 tripods, iron, diam. of ring 10 cm. 

10 troughs, gas-collecting — unless sinks are adapted to this 
purpose 

10 m. tubing, rubber, for Bunsen burners, inside diam. 6 mm. 
7 m. tubing, rubber, soft, inside diam. 5 min'. 

2 voltmeters, student type 

10 wing tops for Bunsen burners 

GLASSWARE 


24 beakers, 100 cc. 

24 beakers, 150 cc. 

24 beakers, 250 ce. 

20 beakers, 400 cc. 

12 beakers, 600 cc. 

12 boats, porcelain, small 
50 bottles, wide-mouth, 250 cc. 
12 bottles, wide-mouth, 500 cc. 
10 burettes, 50 cc. 

15 flasks, Erlenmeyer, 250 cc. 
15 flasks, Florence, 100 cc. 

15 flasks, Florence, 500 cc. 

12 funnels, 60°, 6.5 cm. 

12 funnels, 60°, 9 cm. 

12 graduates, 10 cc. or 25 cc. 


LIST OF SUPPLIES 


viii 

12 graduates, 100 cc. 

12 pipettes, 10 cc. 

50 plates, glass, 8 cm. X 8 cm. 

12 plates, cobalt glass, 8 cm. X 8 cm. 

0.5 kg. rods, glass, diam. 4 mm. 

12 tubes, calcium chloride, 15 cm. 

12 tubes, test, hard glass, 15 cm. X 15 mm. 

12 tubes, test, soft glass, 18 cm. X 25 mm. 

15 doz. tubes, test, soft glass, 15 cm. X 20 mm. 

15 tubes, thistle, 25 cm. stem, 5 mm. diam. 

10 tubes, U, diam. 1 cm. 

4 m. tubing, hard glass, combustion, diam. 15 mm. 

2 kg. tubing, soft glass, medium walls, outside diam. 6 mm. 
24 watch glasses, diam. 8 cm. 

This list does not include reagent bottles for desks and side 
shelf. The size and number of such bottles will be dependent 
upon the number of students taking the course. 

CHEMICALS 

0.5 kg. acid, acetic, 30% 

100 g. acid, fluosilicic 
5 kg. acid hydrochloric, sp.gr. 1.19 
5 kg. acid, nitric, sp.gr. 1.42 
100 g. acid, perchloric 
5 kg. acid, sulfuric, sp.gr. 1.84 
50 g. acid, tartaric 
2 1. alcohol, ethyl, 95% 

25 g. agar, flake 
100 g. aluminum, sheet or wire 
100 g. aluminum chloride 
0.5 kg. aluminum sulfate, cryst 
200 g. ammonium carbonate 
0.5 kg. ammonium chloride 
5 kg. ammonium hydroxide, sp.gr. 0.9 
50 g. ammonium molybdate 
100 g. ammonium nitrate 
100 g. ammonium oxalate 
100 g. ammonium sulfate 
50 g. antimony, powd. 


LIST OF SUPPLIES 


IX 


100 g. antimony trichloride 
100 g. arsenic trioxide 
200 g. barium chloride 
100 g. barium nitrate 
25 g. bismuth 
50 g. bismuth trichloride 
100 g. bromine 
50 g. cadmium nitrate 
2 kg. calcium carbonate, marble chips 
100 g. calcium magnesium carbonate, dolomite 
0.5 kg. calcium chloride, anhydrous, granular 
100 g. calcium fluoride, powd. 

100 g. calcium nitrate 

1 kg. calcium oxide, quicklime, in tin can 
0.5 kg. calcium phosphate, bone ash 
0.5 kg. calcium sulfate, gypsum, powd. 

1 kg. calcium sulfate, plaster of Paris, fine 
0.5 kg. carbon disulfide 
0.5 kg. charcoal, animal 
100 g. charcoal, wood, powd 
200 g. chromic anhydride 
100 g. chromium nitrate 
100 g. cobalt nitrate 
0.5 kg. copper sheet, 1 mm. thick 
1 kg. copper turnings 
1 spool copper wire #24 
100 g. cupric chloride 
100 g. cupric nitrate 
100 g. cupric oxide, powd. 

1 kg. cupric sulfate, blue vitriol 
0.5 kg. ether 
200 g. glucose 

200 g. hydrogen peroxide, 3% 

50 g. iodine, resublimed 
200 g. iron filings 
200 g. iron nails, small 
1 spool iron wire, piano 
100 g. iron chloride, ferric 
100 g. iron oxide, ferric 
100 g. iron oxide, magnetic 
200 g. iron sulfate, ferrous 


X 


LIST OF SUPPLIES 


1 kg. iron sulfide, ferrous 
200 g. lead filings 
200 g. lead nitrate 
100 g. lead dioxide 
100 g. lead monoxide, litharge 
100 g. lead tetroxide, red lead 
50 g. magnesium powder 
50 g. magnesium ribbon 
200 g. magnesium carbonate 
200 g. magnesium chloride 
0.5 kg. magnesium sulfate, Epsom salt 
0.5 kg. manganese dioxide, powd. 

100 g. manganese nitrate 
200 g. mercury 
50 g. mercuric chloride 
100 g. mercuric nitrate 
200 g. mercuric oxide 
100 g. mercurous nitrate 
10 g. methyl orange 
100 g. nickel nitrate 
0.5 kg. paraffin, solid 
10 g. phenolphtalein 
50 g. phosphorus, red 
50 g. phosphorus, white 
0.5 kg. potassium aluminum sulfate, alum 
100 g. potassium bromide 
200 g. potassium carbonate 
0.5 kg. potassium chlorate 
200 g. potassium chloride 
0.5 kg. potassium dichromate 
100 g. potassium ferricyanide 
100 g. potassium ferrocyanide 
0.5 kg. potassium hydroxide, sticks 
100 kg. potassium iodide 
0.5 kg. potassium nitrate 
200 g. potassium permanganate 
100 g. potassium sulfate 
200 g. potassium thiocyanate 
100 g. silica, fine, infusorial earth or kieselguhr 
100 g. silver nitrate 
50 g. sodium 


LIST OF SUPPLIES 


xi 


100 g. sodium ammonium phosphate, microcosmic salt 
0.5 kg. sodium bicarbonate, baking soda 
0.5 kg. sodium carbonate, anhydrous 
2 kg. sodium chloride, table salt, fine 
0.5 kg. sodium chloride, rock salt, coarse 
200 g. sodium dichromate 
1 kg. sodium hydroxide 
0.5 kg. sodium nitrate 
200 g. sodium nitrite 
100 g. sodium peroxide 
200 g. sodium phosphate, secondary 
200 g. sodium potassium tartrate, Rochelle salt 
0.5 kg. sodium sulfate, Glauber's salt 
200 g. sodium sulfite 
0.5 kg. sodium tetraborate, borax 
200 g. sodium thiosulfate 
200 g. starch 
100 g. strontium nitrate 
0.5 kg. sugar, cane, granulated 
0.5 kg. sulfur, flowers 
0.5 kg. sulfur, roll 
200 g. tin, granulated 
100 g. tin chloride, stannic 
100 g. tin chloride, stannous 
1 kg. zinc, granulated 
0.5 kg. zinc, sheet 
100 g. zinc nitrate 
50 g. zinc oxide 
100 g. zinc sulfate 
0.5 kg. Wood's metal 

A very few articles of common household use, such as lard, 
calico, wool yarn, and baking powder are not included in the 
above list. 


TO THE TEACHER 


This manual provides a sufficient number of experiments for 
one year of General Chemistry laboratory, including an intro¬ 
duction to qualitative analysis. Ordinarily the exercises in 
Part I will be performed during the first semester, those in 
Part II during the second semester. 

If only one semester of laboratory work is given in General 
Chemistry , it is suggested that certain supplementary exercises 
may be chosen from Part II. 

In courses designed for students who enter college with 
credit in chemistry, it may be advisable to omit some of the 
simpler exercises of Part I and select further experiments of 
general interest from Part II. 


LABORATORY EXERCISES 

In General Chemistry 
PART I 

THE NON-METALS 

GENERAL LABORATORY RULES 

The following points are set down for the guidance of the 
student. Each item is important and must be strictly observed. 

1. Apparatus and surroundings must be kept clean. 

2. Return bottles of reagents to their proper places on desk 
and side shelf immediately after using. 

3. Never return any material to the reagent bottles. 

4. Do not use more material than necessary. Chemicals are 
expensive and results are often obscured by an excess. 

5. Do not discard anything you have obtained until you are 
sure you have no further use for it. Throw solids into waste 
jars and liquids into the sink. Acids should be washed out 
of pipes with a large amount of water. 

6. Read the label carefully before taking anything from a 
bottle. Using the wrong material may result in serious 
injury. Never carry side shelf bottles to your desk. Use 
clean test tubes for liquids and small squares of paper for 
solids unless otherwise directed. 

7. When heating a substance in a test tube be careful not to 
aim it at anyone who is near. 

8. Exercise particular care in handling hot apparatus, strong 
acids and alkalies, inflammable substances (warning will 
be given), and glass tubing. A glass tube should be held in 
a towel when it is being pushed through a stopper. In case 
of personal injury, however slight, consult an instructor at 
once. 

9. Study the assignment before coming to the laboratory and 
plan your work for best efficiency. Unless this is done the 
required amount of work can not be done in the laboratory 

1 


2 


GENERAL CHEMISTRY 


period. It is often possible to perform two experiments 
simultaneously. Be prepared to discuss with your in¬ 
structor what you are going to do and why. 

10. Laboratory notes on an experiment must be written in your 
notebook as you perform that experiment. The notebook 
is not to be taken from the laboratory except by special 
permission. 

11. Your work must be performed independently, but do not 
hesitate to ask your instructor for advice. Use your text¬ 
book for reference. Above all, be awake — think! 

12. Just before leaving the laboratory make sure that the gas 
and water are turned off, that all bottles are in place, that 
your desk top is clean. 


SUGGESTIONS FOR WRITING LABORATORY NOTES 

1. Carefully note the title and subtitle of the exercise, realizing 
that each experiment is intended to illustrate some impor¬ 
tant chemical fact or principle. Forget for the moment that 
this experiment has previously been performed by others 
and feel that for you it is an original scientific research — a 
“question put to nature.” With this point of view you will 
see the importance of making accurate observations and 
careful conclusions. 

2. Record what you do without repeating exact directions; a 
simple drawing may save a description of the apparatus. 
Observe and record immediately any changes that occur. 
Write out your conclusions from the phenomena observed. 
When any chemical change is involved write the equation for 
the reaction. Answer in your write-up all questions either 
direct or implied by (?). In all be brief, concise and to the 
point of the experiment. 

3. Since these are notes and not literary essays, emphasis is 
placed upon the accuracy and detail of your record rather 
than upon its form. Use the left-hand side of the page for 
drawings and calculations if you so desire. If part or all of 
an experiment is unsatisfactory, mark “Repeated” across 
the old record so it will not be graded and rewrite that por¬ 
tion. Neatness here is of secondary importance to clarity; 
yet a reasonable degree of neatness is required in your notes 
and experimentation. 


LABORATORY EXERCISES 


3 


4. Notes or data, such as weighings, recorded on scraps of 
paper will be taken up by the instructor and destroyed. 
Make all records in your notebook directly and immediately. 
It is both unscientific and inefficient to delay or transfer 
records of an observation. Whenever possible arrange your 
data in the form of a table so you may easily draw com¬ 
parisons and for the convenience of your instructor. 

5. Use your textbook for reference constantly in performing 
and writing up laboratory experiments. Textbook ref¬ 
erences are given in this manual to General Chemistry 
by B. Smith Hopkins, D. C. Heath and Company, 1930. 
Read and follow the “Suggestions on How to Study Chem¬ 
istry,” Textbook: vii and viii. 

6. A FORMAL REPORT will be required on certain quan¬ 
titative experiments. For such experiment a record of data 
is made in the notebook as usual and a copy of this is taken 
home where a more elaborate write-up is prepared. This is 
to be enclosed in a manila paper cover and handed in before 
a date announced by the instructor. In this formal report 
should be included a complete statement of the purpose and 
method of procedure in the experiment, a drawing of the 
apparatus, results (data in tables), calculations and con¬ 
clusions. Neatness and arrangement as well as accuracy 
are important features of this report. 

PROCEDURE FOR THE FIRST LABORATORY PERIOD 

1. Check the apparatus in the desk assigned by your instruc¬ 
tor. You should learn at this time the names of any un¬ 
familiar pieces of apparatus. If anything is missing note it 
on the card, have this countersigned by your instructor, and 
obtain the articles at the storeroom. When your outfit is 
complete sign the fist and return it to the storeroom. 

2. Read carefully the General Laboratory Rules. Read 
them again at home and as often as necessary. 

3. Proceed with Exercise 1. Practice working glass tubing 
according to the instructions and when sufficient technique 
has been acquired prepare a wash bottle as in Figure 4. 
Submit this to your instructor and, if approved, fill it with 
distilled water and keep it in your desk for use when 
directed. No record of this exercise is required. 


EXERCISE 1 


THE USE OF GLASS TUBING 



^ )\ 


Fig. 1 


(а) To Cut Glass Tubing. — By means of a triangular file 
make a transverse scratch upon the glass tubing at the desired 
point. Hold the tube so that the ends of the thumbs are to¬ 
gether on the tube opposite the 
scratch as in Figure 1. Press out¬ 
ward with the thumbs and pull 
gently in the opposite direction 
with both hands. To cut wider 
glass tubing a deep scratch should 
be made completely around the 
tubing and a crack started by 

touching the scratch with the red-hot end of a glass rod. 

(б) To Fire-Polish Glass. — Always round off the edges of 
freshly cut glass tubing or glass rod in order to avoid injury to 
your hands and to prevent cutting the rubber stopper or rubber 
tubing into which the glass may be inserted. Fire-polishing is 
accomplished by rotating the end of the glass tube or rod in a 
Bunsen burner flame until the sharp edges become smooth. 
Why does this occur? If this treatment is continued too long 
with small glass tubing the end of the tube will close com¬ 
pletely. 

(c) To Bend Glass Tube. —Attach a “wing-top” to a 
Bunsen burner, close the air holes at the bottom, and light 
the burner. Hold a piece of /t^, K fXTr r? 
tubing lengthwise in the yellow ■ v ' : - 

luminous flame, as in Figure 2, 
not across the flame, rotating the 
tubing slowly to insure uniform 
heating. Continue heating the 
tube in this manner until it is 
soft enough to bend of its own 
weight, then remove it from the flame and at once bend it 
to the desired angle. Try to produce a bend that will coincide 

4 


V 




Fig. 2 













LABORATORY EXERCISES 


5 



with some angle drawn on a piece of paper. The hot glass 
should not come in contact with the varnished surface of the 
desk but should be suspended in the air or laid upon a piece of 
asbestos to cool. 

The bends A and B in Figure 3 are correctly made; C and D 
are unsatisfactory because of the constriction at the bend and 
because they are more likely to break. 

Bends such as C and D result from non- 
uniform heating of the glass, heating the 
tube through too short a length, bend¬ 
ing the tube while it is in the flame, or 
forcing the glass to bend before it has 
been heated to the required softness. 

(d) To Make a Jet. — Remove the 
wing-top, light the burner, and open the 
air holes at the bottom to produce a 
hot, non-luminous flame. Heat a piece 
of glass tubing at a distance of 2 or 3 
inches from one end, rotating the tubing 
in the hot flame until the glass begins to 
collapse. Remove the tubing from the 
flame and pull the ends apart in a 
straight line until the capillary has the desired size. When the 
glass is cool, carefully cut the capillary and fire-polish both 
ends of the jet. 

( e ) To Insert a Glass Tube in a Stop¬ 
per. — Beginners in the chemical laboratory 
probably suffer more injuries from careless¬ 
ness in this simple operation than from any 
other single cause. There is no danger from 
this source if the following precautions are 
observed: See that both ends of the tube are 
properly fire-polished. Be sure that the tube 
is cold, and wet both tube and stopper. Push 
the tube through the hole in the stopper 
with a screw-like motion. If the tube is 
fragile or has a bend in it, protect your 
hand with a towel. Do not exert a strain upon a bend in 
the tube. 

(/) Construction of a Wash Bottle. — Obtain a two-hole 
rubber stopper from the storeroom to fit the 500-cc. Florence 




Fig. 4 
















6 


GENERAL CHEMISTRY 


flask, prepare the other parts, and assemble a wash bottle as 
in Figure 4. 

Before coming to the next laboratory class read carefully 
Suggestions for Writing Laboratory Notes and study 
Exercise 2. 


EXERCISE 2 

MEASUREMENT — THE METRIC SYSTEM 

Textbook: 5^-55 

The metric system of measurement is used for most scientific 
work and is gradually coming into general use. It has three 
fundamental units — the meter, the liter, the gram. These 
units are interrelated and all multiples and submultiples are 
decimal. Refer to table near back of this book. 

(a) Length. — Examine a meter stick on both sides. Cal¬ 
culate the percentage (of a yard) by which a meter exceeds a 
yard. Measure a length of more than two feet in both inches 
and centimeters and from these values calculate the number of 
centimeters in an inch. 

( b ) Volume. — A liter is the volume of a cube having edges 
1/10 (one-tenth) meter in length. Draw a liter cube and label 
it to show the relation of the length of edge to volume. How 
many cubic centimeters (cc.) are there in a liter? 

Measure the volume of an ordinary test tube by means of the 
graduated cylinder. Mark with gummed labels the test tube 
levels that correspond to 10 cc. and 20 cc. This tube may be 
used in your work for approximate measurements. 

(c) Weight and Density. — Weigh an empty beaker on the 
laboratory scales. Fill it about half-full of water and weigh 
again. Measure the water and calculate the weight of 1 cc. of 
water and of 1 liter of water. Label the cube drawn above with 
the weight of water required to fill it. Your drawing now dem¬ 
onstrates that, in the metric system, the units of length, volume 
and weights are related. 

If the density of concentrated sulfuric acid is 1.8 (density of 
water is 1), what is the weight of the contents of the bottle of 
that acid on your desk? Its volume is one-fourth liter. 

Determine the density of some solid as follows: Carefully 
weigh about 20 g. of some insoluble material — e.g., broken 
glass tubing, marble chips, ferrous sulfide, or iron nails. Fill 


LABORATORY EXERCISES 


7 


the 100-cc. graduate with water to a definite mark (read oppo¬ 
site the bottom of the curved water surface or meniscus), add 
the weighed material and note the increase in volume. From 
this data calculate the density (specific gravity) of the material 
to the second decimal place. State in words and by a mathe¬ 
matical formula the relationship of density, weight and volume. 

Complete the following by supplying the correct number: 


0.8 meter 
2364 millimeters 
1.75 grams 
1250 cc. 

2.4 kg. of water 


centimeters 

meters 

milligrams 

liters 

cubic centimeters 



5% Difference 



10% Difference 


5uBO«/.oeo : 3 • 36 • S7fe 

Yard 

S’JBOIVIOC.D •_1 0 -100 - IOOO 

Meter 

10% Difference 

Fig. 5 


LABORATORY TECHNIQUE 

It is very important to form correct habits in the laboratory. 
The best method for performing any operation should be learned 
by the student when that operation is first encountered, and this 
method should be used thereafter on all occasions. A few gen¬ 
eral instructions were given under General Laboratory Rules. 
Refer to these from time to time until you are sure you are 
observing all of them as a matter of habit. Further definite 
instructions for correct procedures will be presented in the exer¬ 
cises as they are needed; the following processes of manipula¬ 
tion should be learned at this point for immediate use. 

(a) Removing Solids from Bottles. — The required amount of 
a solid should be removed from the bottle by means of the hollow 
glass stopper as indicated in Figure 6. If the solid is to be 
weighed on the laboratory scales, two pieces of paper of equal 












8 


GENERAL CHEMISTRY 


weight are placed, one on each platform of the scales, the weights 
are placed on the right-hand platform or the rider is moved to 
the right, and the solid is then slowly poured upon the paper on 

the left-hand platform by 
tilting and gently rotating 
the stopper until the scales 
are balanced. 

(6) Pouring Liquids 
from Bottles.—When 
using a reagent from a 
glass-stoppered bottle, re¬ 
move the stopper by tak¬ 
ing it between the fingers 
at the back of the right 
hand as shown in Figure 7a. 
The presence of the stop¬ 
per does not interfere with 
the use of that hand for 
holding the bottle. Never lay the stopper down as it will 
become contaminated and may, when returned to the bottle, 
make the contents unfit for use. In pouring from a bottle 
the neck of the bottle should touch the edge of the receiving 
vessel (Fig. 7b). This prevents the liquid from running back 



Fig. 7 



on the side of the bottle and is especially important in the 
handling of strong acids or other corrosive liquids. 

Never dip a stirring rod, pipette, platinum wire or other 
apparatus into a reagent bottle. To obtain a drop of the re¬ 
agent on a stirring rod, pour some of the reagent into a watch 
glass or test tube and dip the rod into this. 





















LABORATORY EXERCISES 


9 


(c) Filtering. — Prepare the filter paper by folding it along 
its diameter and then folding it again at a right angle to the first 
fold (Fig. 8). Open the paper to form a cone, fit this in a 
funnel, and wet the paper with a few cubic centimeters of dis¬ 
tilled water to cause 
it to stay in place. 

Pour out any excess 
water before using the 
filter. 

Fig. 9 shows the cor¬ 
rect arrangement for 
filtering. The liquid 
should be poured 
slowly along the stir¬ 
ring rod and the filter 
paper should not be- p IG< g p IG> 9 

come too full. The 

filtrate, passing through the stem of the funnel, should run down 
the side of the receiving vessel and should not be allowed to 
fall in drops because this would cause spattering. 

If the precipitate is very finely divided it may run through the 
filter paper at first. If the first portion of the cloudy filtrate is 
again filtered it will probably come through clear since the filter 
paper becomes more effective after some of the precipitate has 
been collected in the paper. 

If a precipitate is heavy and settles well, most of the liquid 
may be separated from it by carefully pouring the liquid away 
from the solid. This process is called decantation. For a more 
complete separation the liquid may be decanted through a 
filter. 

(, d ) Heating Materials in Glass and Porcelain. — Thick- 
walled vessels of glass or porcelain must never be heated. When 
heating apparatus of glass or porcelain remember to raise the 
temperature gradually. Sudden heating or cooling may cause 
the apparatus to crack. 

Beakers, flasks and evaporating dishes should be heated on a 
wire gauze and not in contact with the direct flame. 

When heating a liquid in a test tube, raise the temperature 
gradually and uniformly throughout the depth of the liquid to 
avoid the sudden formation of vapor below the surface of the 
liquid. Uniform heating may be accomplished by passing the 












10 


GENERAL CHEMISTRY 


tube back and forth through the flame and agitating the con¬ 
tents of the tube, or by waving the flame along that part of the 
tube occupied by the liquid. 

EXERCISE 3 

PROPERTIES OF MATTER — CHEMICAL AND 
PHYSICAL CHANGE 

Textbook: 18-22 

(а) Examine a clean piece of copper wire or copper foil and 
note its properties. Hold the piece of copper by means of a pair 
of forceps in the outer flame of the Bunsen burner until the 
copper is red-hot. Remove the copper from the flame and 
observe it while it cools in the air (?). Compare the properties 
of the product upon the surface of the copper with the proper¬ 
ties of the original copper. 

Repeat the experiment with a platinum wire and with a piece 
of magnesium ribbon. Tabulate the properties of each metal 
before and after heating. State in each case whether a chemical 
or physical change was brought about by the application of 
heat to the metal. 

Place the piece of platinum wire, a clean piece of copper and 
a piece of magnesium ribbon in separate test tubes and add 3 
or 4 cc. of dilute hydrochloric acid to each. Is there evidence 
of chemical change in any of these? Repeat this part of the 
experiment using a few drops of concentrated nitric acid with 
each metal. Describe the results and state the evidence for 
any chemical changes observed. Wash the acid from the piece 
of platinum wire and preserve the wire for subsequent use. 
Summarize the results you have obtained to show the differ¬ 
ences in the chemical properties of the three metals. 

Define briefly and illustrate: physical property, chemical 
property, physical change, chemical change. 

(б) Obtain about 2 g. of sodium carbonate, note its properties 
(color, taste, etc.) and dissolve it in 30 cc. of water, warming 
the water to hasten solution. Divide this solution into three 
equal portions and proceed with the separate portions as follows: 

1. Evaporate this portion to dryness in a porcelain dish and 
compare the residue with the original substance. 

2. Add dilute hydrochloric acid, a few drops at a time with 


LABORATORY EXERCISES 


11 


stirring, to this portion until no further action is observed and 
evaporate this portion to dryness in a porcelain dish. Examine 
this residue and compare it with the original solid with respect 
to appearance and taste. 

3. To the third portion add a solution of calcium chloride 
until no further change is observed. Filter out the precipitate, 
transfer it to a small beaker and dissolve it in the minimum 
quantity of dilute hydrochloric acid (?). Evaporate this solu¬ 
tion to dryness and examine the residue. It is calcium chloride. 
State, in the order of their occurrence, the changes which have 
been brought about in this portion of the sodium carbonate 
solution. If the filtrate obtained in this experiment had been 
evaporated to dryness, what residue would have been obtained? 

Record and learn the formulas for all of the substances used 
in this exercise. Write the equations for each of the chemical 
changes observed in this exercise. 

EXERCISE 4 

ELEMENTS, COMPOUNDS — SEPARATION FROM 
MIXTURES 

Textbook: 2J+-28 

The substances to be found in the laboratory, if pure, may 
be classified as elements and compounds. Define each class. 

Write in your laboratory notebook a brief description of four 
elements found on the side shelf. Two should be metals and 
two non-metals. Give name and symbol and such properties 
as color, form, hardness, etc. 

Find two soluble compounds and two insoluble compounds. 
Give names and formulas and describe briefly. To test solu¬ 
bility take a very small quantity in one-half test tube of water 
and shake — heat if necessary. A solution must be clear and 
transparent if not too deeply colored. State the distinction 
between symbol and formula. 

Now place about a gram of a soluble substance and about the 
same quantity of an insoluble substance (both finely divided) in 
a small beaker, add 20 cc. of water, heat to boiling and decant 
through a filter. What is the residue? Evaporate the filtrate 
to a small volume and allow to cool. Identify the crystals. 
Has a chemical change occurred? 


12 


GENERAL CHEMISTRY 


Two soluble substances may be separated by fractional crys¬ 
tallization: in general, the less soluble one will crystallize first. 
Referring to the solubility curves in your textbook, is sodium 
chloride or potassium nitrate less soluble in cold water? In 
boiling water? Which curve is most influenced by temperature 
change? 

Take 3 g. common salt (sodium chloride, formula?) and 10 g. 
saltpeter (potassium nitrate, formula?) in a test tube, add 15 cc. 
of water; warm until the salts are dissolved. Cool the solution 
by placing the tube in a beaker of cold water. Note the shape 
of the crystals; taste them. What are they? Pour off the 
mother liquid into a small beaker and evaporate it slowly until 
the crystals appear in the hot liquid. Note their form and 
taste. What are they? As the remaining solution now cools 
what crystals appear? 

Briefly indicate how you could separate two mutually soluble- 
liquids such as alcohol and water. 


EXERCISE 5 

STUDY OF A MIXTURE 

(a) Gunpowder is a mixture of saltpeter (potassium nitrate), 
sulfur and charcoal (carbon). Grind separately to a fine 
powder 3 g. saltpeter, 0.5 g. charcoal, and 
0.5 g. sulfur. Mix these JlJoroughly on a 
piece of paper. Place half the mixture in a 
dry test tube, add 5 cc. of carbon disulfide, 
shake well (closing test tube with thumb) 
and filter,* bringing all the solid residue on 
the filter paper. ( Caution: Carbon disulfide 
is very inflammable; do not bring it near a 
flame.) Catch the filtrate in a watch glass 
and allow it to evaporate spontaneously. 

Allow the residue in the filter paper to 
dry, then pour 10 cc. of warm water through 
it, a little at a time. Catch this filtrate in a porcelain dish and 
evaporate it almost to dryness on a wftfrer - bath (Fig. 10). 
Identify the crystals. What is left on the watch glass? On 

• Do not wet the filter paper with water in this case. 



Fig. 10 













LABORATORY EXERCISES 13 

the filter paper? Carbon is not soluble in any laboratory 
reagent. 

(b) Make a conical pile of the remainder of the gunpowder on 
an asbestos mat and ignite it by means of a paper fuse or the 
Bunsen burner. Compare the volume of the residue with that 
of the gunpowder used. To what form was most of the material 
converted? List the properties of gunpowder which you think 
make it useful as an explosive or propellant. 


EXERCISE 6 

CHEMICAL CHANGE 

(a) Combination of Two Elements. — Grind together 7 g. 
fine iron filings and ^gf sulfur. Examine some of the mixture 
with a lens. Is it homogeneous? Spread some of the mixture 
on a piece of paper, draw a magnet under it or through the 
mixture and blow away the sulfur (?). State briefly another 
method by which you could separate the iron and sulfur. ^ (Re¬ 
fer to previous exercise.) 

Save out a small portion of the mixture for later comparison, 
put the remainder in a dry test tube, and while holding the tube 
in an inclined position apply a vigorous flame to the lower end 
until the mixture begins to glow. Remove the tube from the 
flame and notice the progression of the glow through the mass. 
What is the source of this heat? One characteristic of a chem¬ 
ical change is the transformation of chemical energy into heat, 
or vice versa. 

When the tube is cool, break it and examine the product, 
comparing it with the original mixture. Is it homogeneous? ’ 
Magnetic? Can the two elements now be separated by means-^.^ 
of CS 2 ? ( What has taken place?) In test tubes add a little 
dilute hydrochloric acid to some of this product and to some 
of the original mixture. Determine by odor whether the same 
v 3^as is evolved in both cases. A second characteristic of chem¬ 
ical change is the formation of one or more new substances. 

If all the mixture had been heated, what weight of iron sul¬ 
fide would have been formed? The third characteristic of 
chemical change is constancy of total weight. 

What is the formula of a compound of the above proportions? 


GENERAL CHEMISTRY 




14 

Refer to relative (atomic) weights of iron and sulfur. Cal¬ 
culate also the formula of a compound consisting of seven 
parts iron to eight parts sulfur. See Model Example II. 

(6) Decomposition of a Compound. — Prepare a tube about 
10 cm. long by taking a longer piece of glass tubing, heating it 
10 cm. from one end, and drawing the two pieces apart until a 
closed tube is madey Put a small quantity of mercuric oxide 
(formula?) into this tube and heat it. Insert a glowing splinter 
of wood into the tube (?). This is a test for oxygen. What 
other substance is formed? Note down a few of its properties. 
Restate the results of this experiment in terms of the three 
characteristics of chemical change, 

What change in color is noticed as the mercuric oxide is 
heated and as it cools? Is this, a physical or chemical change? 


EXERCISE 7. 

SOURCES OF OXYGEN 

Textbook: 30-33 

(а) Air is a mixture of gases, about one-fifth being oxygen. 
How may oxygen be separated from this mixture? How is 
oxygen obtained from water? 

(б) To prepare oxygen in the laboratory some oxygen com¬ 
pound is usually decomposed by heat as in Exercise 6 ( b ). 
Write the formulas for the following substances and obtain 
samples (about 0.5 g.) from the shelf: ferric oxide; potassium 
chlorate; sodium carbonate; manganese dioxide; potassium 
nitrate. Heat each of these separately in tubes prepared as in 
Exercise 6 (5), or in the small hard glass test tube, and test for 
oxygen (?). Do you know any other compounds that yield 
oxygen on heating? What would you call that property which 
determines whether or not an oxygen compound will liberate 
oxygen when heated? 

( c ) Place one-half gram sodium peroxide (storeroom) in a 
test tube, add a few drops of water and test for the evolution 
of oxygen (?) Add more water and test the solution with red 
litmus paper (?). Write a word equation to show the react¬ 
ants and products in this change. Repeat the equation using 
formulas. 



LABORATORY EXERCISES 


15 


EXERCISE 8 

CATALYSIS 

Textbook: 33-34 

(a) Place a gram of potassium chlorate in a dry test tube, 
clamp this on the iron stand in an inclined position and heat 
carefully with a Bunsen burner until the salt melts. Have 
ready at hand a small amount of manganese dioxide. Heat 
the potassium chlorate slowly, just above its melting point, so 
that bubbles of gas rise very slowly through the liquid. Test 
for oxygen (?). Remove the flame, add the manganese dioxide 
and note the effects Test again with a glowing splinter (?). r 

(b) Heat the mixture until the potassium chlorate is entirely 
decomposed. Test (?). Allow it to cool, add 10 cc. of water, 
shake to disintegrate the residue and pour through a filter paper. 
Catch the filtrate in a porcelain dish and evaporate it to dry¬ 
ness. Compare it by taste and appearance with potassium 
chlorate. Does it contain oxygen? How can you prove it? 

What is the residue on the filter? How could you prove its 
identity? Explain its use in this experiment. Can manganese 
dioxide be made to give up a part of its oxygen (Exercise 7)? 
Is ferric oxide decomposed by heating (Exercise 7)? 

(c) Repeat (a) using ferric oxide in place of manganese di¬ 
oxide. 


EXERCISE 9 

PREPARATION OF OXYGEN; COMBUSTION — 
FORMATION OF ACIDS AND BASES 

Textbook: 33-36 

Make a mixture of 10 g. of potassium chlorate and 3 g. man¬ 
ganese dioxide. Place this in a test tube provided with a one- 
hole stopper and a delivery tube. Arrange the apparatus as 
shown in Figure 11. The lower end of the delivery tube must be 
curved upward so a bottle may be inverted over it. The test 
tube should not be more than half full of the mixture. Fill four 
bottles with water, cover with glass plates and invert them in the 
sink or gas-collecting trough which is partly filled with water. 


16 


GENERAL CHEMISTRY 


Place a bottle filled with water over the end of the delivery 
tube and gently heat the mixture in the test tube. Do not 
allow the molten contents of the test tube to come in contact 
with the stopper. Collect four bottles of oxygen in this manner. 
As each bottle is filled, slip a glass plate over its mouth and 
set the bottle upright on the desk for use in the following 
experiments: 

(a) Carbon. — Wrap a piece of iron wire around a bit of char¬ 
coal or place the charcoal in a clean iron deflagrating spoon. 
Heat this in the Bunsen burner until the charcoal glows and then 
lower it into a bottle of oxygen (?). Do not drop the charcoal 
into the bottle. When the charcoal no longer burns, remove 
it and add 10 cc. of water, shake well and test with blue lit¬ 
mus (?). Pour part of the solution into a test tube and add 
about 5 cc. of limewater (calcium hydroxide) (?). The pre¬ 
cipitate is calcium carbonate, having the same composition as 
chalk, marble and limestone. 

Put 5 cc. of limewater in a test tube and blow your breath 
through it by means of a glass tube (?). For what gas is lime- 
water a test? 

Write equations for the combustion of charcoal; for the addi¬ 
tion of water to the product; for the addition of limewater. 

(i b ) Sulfur. — Burn some sulfur in the iron spoon in another 

bottle of oxygen (?). No¬ 
tice the odor. Add water, 
shake and test with blue lit¬ 
mus (?). Test the dilute 
acids on your desk with lit¬ 
mus. What element is com¬ 
mon to all acids? Name 
two acids found in foods or 
beverages. 

(c) Phosphorus. — ( Cau¬ 
tion: White phosphorus is 
kept under water because it 
burns spontaneously in the 
air. It must be handled 
with forceps. If it takes 
fire in contact with the flesh it produces very painful burns.) 

In a dish of water obtain a small piece of phosphorus from 
the storeroom. Lay it for a moment on a piece of filter paper 












LABORATORY EXERCISES 


17 


to dry it, then place it in the deflagrating spoon. Ignite the 
phosphorus by touching it with a slightly warm file or piece of 
glass tubing and at once lower it into a bottle of oxygen. What 
is the appearance of the product? Add water and test with 
litmus (?). 

(d) Magnesium. — Obtain a piece of magnesium ribbon 
about 5 cm. long. Hold the ribbon in forceps, ignite.it in the 
Bunsen flame and thrust it into a bottle of oxygen. The light of 
burning magnesium is very rich in ultra-violet rays. What 
use is made of magnesium in photography? Examine the 
product (?). Add water and test with red and blue litmus 
papers (?). Classify as metals and non-metals the four ele¬ 
ments which you burned in oxygen. State the reactions by 
which: 




1. An acid may be formed from a non-metal. 

2 . A base may be formed from a metal. 


In order that these two statements may be true the oxide must, 
of course, be soluble in water. 

Be sure you have written equations for all the chemical 
changes which you observed in this exercise. 


EXERCISE 10 

ACCURATE WEIGHING 

The instructor will explain the operation of the chemical bal¬ 
ance and give directions for its use. Note carefully the follow¬ 
ing points: 

1. The balance is a delicate and expensive instrument; handle 
it with care. If it does not seem to work properly call your 
instructor; do not attempt to adjust it yourself. 

2. The object to be weighed is placed on the pan at the left 
of the observer. Begin with larger weights and work down 
to the correct weighing. 

3. Handle weights with forceps only. Do not borrow weights 
or rider from any other balance. 

4. Never put anything on the pans or remove anything from 
them while the beam is unsupported. Do not release the 
beam very far unless the pans are very nearly balanced. 


18 


GENERAL CHEMISTRY 


5. Chemicals are not to be placed directly on the balance pan: 
use glass or porcelain containers. 

6 . The balance case should be closed while determining the 
zero point or taking the final reading of a weighing. 

7. In an experiment where several weighings are to be made, 
the same balance should be used for all. 

8 . Count weights carefully and check by noting vacant places 
where weights are kept. The weighings must be recorded 
directly in your notebook — never carry scraps of paper 
into the balance room. 

9. After a weighing is completed, always return weights to 
their proper places and raise the rider on its support. See 
that the balance case is left in order and its door closed. 

10. Observe strictly the directions posted in the balance room. 

The instructor will give you an accurately weighed piece of 
metal for practice weighing. Determine the weight of this to 
the fourth decimal place (tenth of a milligram); record it with 
the number of the metal slug and the zero point of the balance 
in your notebook and submit this record to your instructor. 

The Zero Point. — Before making any weighing the “zero 
point” of the balance must be determined. This is the true 
center of the swings which the pointer makes on the lower scale. 
If the balance were in perfect adjustment this zero point would 
coincide with the center of the scale. This will rarely be the 
case as the adjustment is affected by many influences. The 
zero point is found by recording an uneven number of consecu¬ 
tive swings to the left and right. The first and last swings are 
thus on the same side and the shortening of the swings due to 
friction is, therefore, neutralized. The average swing for each 
side is calculated, the difference between these, divided by 2, is 
the zero point. An example should make this clear: 



Left 

Right 

Swing — 

(1) 5.4 

(2) 4.2 


(3) 5.0 

(4) 3.8 


(5) 4.6 


Average — 

5.0 

4.0 


Zero point = 5 (left) minus 4 (right) -f- 2 = 0.5 to the left. 
The true center of the swings is hence one-half division on the 
ivory scale to the left of the center line. 




LABORATORY EXERCISES 


19 


Weighing. — In making the weighing the weights aie ad¬ 
justed until the pointer is swinging equally on both sides of the 
zero point just found. This may be accurately determined by 
recording the swings and making a calculation as for the zero 
point. The zero point of the loaded balance should closely 
coincide with the zero point when the pans are empty. 

The most common error of the beginner is the incorrect 
expression or addition of the weights. A weighing should be 
expressed decimally in terms of grams. Thus, the following 
weights: 5 g., 200 mg., 50 mg., 10 mg., rider on 6.4 mg., are 
recorded as 5.2664 g. 


EXERCISE 11 

WEIGHT OF A LITER OF OXYGEN —LAW OF 
CONSTANT COMPOSITION 

Quantitative — Formal Report Required 

In a clean, dry, hard-glass test tube put a small amount 
(about 0.3 g.) of the manganese dioxide provided in the balance 
room. Weigh accurately, expressing the weight to four decimal 
places. Now add about a gram of potassium chlorate (balance 
room) and weigh the tube again. The difference between these 
two weighings represents the accurate weight of what? Why is 
it not necessary to know the exact weight of the manganese 
dioxide? Mix the materials by shaking, being careful not to 
lose any. 

Arrange the apparatus to collect oxygen by displacement of 
water. Use a 500-cc. bottle for this purpose. Heat the mix¬ 
ture until the potassium chlorate is entirely decomposed i.e., 
no more oxygen is evolved — and disconnect the test tube 
from the delivery tube at once to prevent it from drawing in 
water as it cools. Raise or lower the bottle of oxygen until the 
levels of the water inside and outside are exactly the same. Now 
slip a glass plate under the mouth of the bottle, hold it tightly 
against the bottle and remove the bottle and contents to an 
upright position on the desk. The volume of the oxygen may 
now be measured indirectly by measuring a corresponding 
volume of water. Remove the glass plate and, by pouring 
water from a graduated cylinder, carefully note the amount 


20 


GENERAL CHEMISTRY 


required to fill the bottle. Record this measurement as the 
volume of oxygen collected at room conditions. Take the tem¬ 
perature of the water, read the barometer and correct the pres¬ 
sure for the vapor pressure of water at the observed temperature 
(table in back of book). From this data calculate the volume 
which the oxygen would occupy at standard conditions — 0° 
Centigrade and 760 mm. pressure. 

When the tube is cold weigh it again. What does the loss in 
weight represent? What is the per cent of oxygen in potassium 
chlorate according to your data? According to the formula? 

From the corrected volume and the weight of the oxygen cal¬ 
culate the weight of a liter of oxygen under standard conditions. 
Compare your result with the accepted weight for this gas and 
calculate the percentage error. List the probable sources of 
error in your determination. Is the result affected by the fact 
that the test tube and delivery tube were full of air at the begin¬ 
ning of the experiment and full of oxygen at the end? Why is 
it necessary to correct for the vapor pressure of water? Would 
the observed volume of oxygen have been greater or smaller if 
it had been collected over dry mercury? 

Calculate the weight of potassium chlorate required to give 
a liter of oxygen at standard conditions. 

State the law illustrated by the fact that, with careful work, 
different experimenters obtain the same result for the percent¬ 
age of oxygen in various samples of potassium chlorate. 

Each observation should be recorded directly in your note¬ 
book as soon as made. A table such as the following is suggested. 

Table of Data 


Weighing 1. Weight of tube and Mn0 2 . g 

Weighing 2. Weight of tube and Mn0 2 and KCIO 3 .... g 

Weight of KCIO3 used. g 

Weighing 3. Weight of tube and contents after heating g 

Weight of oxygen expelled. g 

Volume of oxygen collected. cc 

Temperature of the water. °C 

Barometric pressure. mm 

Corrected pressure of the oxygen. mm 

Corrected volume of the oxygen (S. C.)... cc 

(Indicate calculation) 










LABORATORY EXERCISES 


21 



EXERCISE 12 


METHODS OF PREPARING HYDROGEN 


Textbook: 75-81 


(a) Action of Water on Metals. — ( Caution: Sodium is kept 
under kerosene and must not come in contact with moist 
fingers. Handle it with forceps.) Pour 10 cc. of water into a 
beaker and drop in a piece of sodium (storeroom) about the 
size of wheat grain (?). Float a piece of filter paper on the water 
and place on it another small piece of sodium (?). Why is the 
kindling temperature reached in this case and not before? 
Test the solution in the beaker with red litmus (?). Rub some 
of the solution between thumb and forefinger (?). Repeat this 
and the litmus test with some sodium hydroxide from the shelf. 
Write the equation for the action of sodium on water. 

Recall the lecture experiment in which steam was passed over 
hot iron (equation). Compare this and the above reaction with 
regard to the type of products and with regard to the amount 
of hydrogen obtained from 36 g. of water. How else may hy¬ 
drogen be obtained from water? 

(i b ) Action of Acids on Metals. — Place a piece of zinc in each 
of three test tubes and cover one with dilute hydrochloric, one 
with dilute nitric and one with dilute sulfuric acid. Test the 
evolved gas with a lighted match or burning splinter (?). A 
slight explosion in the test tube indicates the presence of hydro¬ 
gen. What salt is left in solution in each case? 

In different test tubes try the action of hydrochloric acid on 
bits of the following metals: aluminum, tin, magnesium, zinc, 
iron, copper. If dilute hydrochloric acid has no effect pour 
it off and add the concentrated acid. Arrange the metals in 
the order of the activity with which they appear to displace 
hydrogen. 

Why should not sodium be used with an acid to generate 
hydrogen? What metal might be used to make a container for 
hydrochloric acid? (See Electromotive Series.) 


22 


GENERAL CHEMISTRY 




EXERCISE 13 

PREPARATION AND PROPERTIES OF HYDROGEN 

Textbook: 86-91 

(а) Preparation. — ( Caution: Hydrogen is explosive when 
mixed with air. Do not bring a flame near the hydrogen gen¬ 
erator.) Fit a flask with a rubber stopper bearing a thistle tube 
and delivery tube as shown in Figure 12. Place R) g. of zinc 
in the flask and add enough water to cover thjB end of the thistle 
tube. Pour 2#'cc. of dilute sulfuric acid down the thistle tube 
and add more when needed. Collect the gas in wide-mouthed 
bottles by displacement of water. If the action is slow, add 
through the thistle tube a few cc. of copper sulfate solution. 
The apparent catalysis in this case is caused by metallic copper 
depositing on the zinc and forming with it an electric couple. 
The first bottle of gas collected will be a mixture of air and 
hydrogen and may be discarded or exploded by bringing its 
mouth to a flame. 

(б) Combustion. — Collect a bottle of pure hydrogen and, 
while holding it mouth downward, apply a lighted candle or 
match attached to a wire or piece of glass tubing (?). Thrust 
the candle further up into the hydrogen (?). Withdraw it 
slowly (?). State your conclusion as to the ability of hydrogen 
to burn and to support combustion. What is the color of the 
hydrogen flame? Does it produce much heat? What product 
do you see formed in the bottle? Equation. 

(c) Lightness. — Hold a bottle of air mouth downward and 
pour a bottle of hydrogen upward into it. Now test both bottles 
with a flame (?). From the weights of a liter of the gases (text¬ 
book) indicate the relative weights of air and hydrogen. What 
use is made of this property of hydrogen? What is the objec¬ 
tion to hydrogen for this use? What substitute is there? 

(d) Explosiveness. — Fill a wide-mouthed bottle about a 
third full of hydrogen, drop the remainder of the water out 
allowing air to enter and wait a moment for the gases to mix. 
Apply a flame to this mixture and compare the result with that 
obtained in ( b ). 

(e) Diffusion. — Obtain a cylinder of unglazed porcelain and 
fit it with a rubber stopper bearing a glass tube about 25 cm. 



I 


LABORATORY EXERCISES 23 

long. Let |he end of this tube dip into a beaker of water and 
bring a bottle of hydrogen down over the porous cylinder. If 
the mouth of the bottle is too small pour the hydrogen upward 



into a large beaker placed over the cylinder (?). After a min¬ 
ute remove the vessel of hydrogen and watch the level of water 
inside the tube (?). In your opinion, why can hydrogen pass 
through the pores of the unglazed porcelain, faster than can air? 

(/) Other Product. — rater about ^cc. of the liquid in 
' 4he~4ftsk - • -and' evaporate this on a water 

bath. Describe the residue. To what class of substance does 
it belong? Write the chemical equation to show how it was 
formed. - - lu / 

(g) Problem. — Calculate the weight of hydrogen possibly 
liberated by the weight of the zinc you used if it had all dis¬ 
solved. Calculate the volume of this hydrogen if a liter weighs 
0.089 g. at standard conditions. About what per cent of this 
volume did you actually utilize in this exercise? 

EXERCISE 14 

HYDROGEN AS A REDUCING AGENT 

Textbook: 77-80 

Pour off the liquid from the hydrogen generator used in the 
previous exercise and if sufficient zinc is left use it for this ex- 


' • 


•^sUmXIjlQ 














24 


GENERAL CHEMISTRY 


ercise, providing more if necessary. Prepare a right-angled 
delivery tube A, set up the apparatus as shown in Figure 13 

and prepare hydrogen 
as in Exercise 13 (a). 

(a) Reduction of 
Copper Oxide. — Place 
about 0.5 g. of copper 
oxide in the middle of 
the ignition tube B. 
When the hydrogen 
generator has been 
working long enough to 
displace air from the 
apparatus, wrap a towel 
around it and light the glass tip C by collecting a test tube full 
of pure hydrogen over it and using this tube of hydrogen as a 
torch. (Recall the method as shown in the lecture or have 
your instructor demonstrate.) Never light a hydrogen gen¬ 
erator by any other method. Hold a cold dry beaker over the 
hydrogen flame and note the deposit (?). Hold a piece of iron 
wire in the flame (?). 

Heat the copper oxide with the Bunsen burner; notice the 
effect upon the hydrogen flame (?), any deposit beyond the 
copper oxide (?) and any change in the appearance of the copper 
oxide (?). Explain the results, giving the equation. 

( b ) Reduction of Iron Oxide. — Repeat the experiment, using 
magnetic oxide of iron in place of copper oxide and make obser¬ 
vations as above. Recall the action of steam on iron, Exer¬ 
cise 12 (a), and explain. What kind of reaction is this? What 
determines its direction? Is the reduction of copper oxide 
reversible? Can copper ever displace hydrogen? Recall 12 (b). 

• 

EXERCISE 15 

WEIGHT OF METAL TO DISPLACE A GRAM ATOM 
OF HYDROGEN 

Quantitative — Formal Report Required 

Obtain a large test tube, fit it with a two-holed stopper and 
set up the apparatus as illustrated in Figure 14. Weigh ac- 



Fig. 13 













LABORATORY EXERCISES 


25 


curately a strip of magnesium ribbon about 6 cm. long. Place 
this strip in the test tube and fill the whole apparatus with 
water by pouring it into the funnel and regulating the flow by 
means of the pinch- 
cock. Be careful to 
drive out any air bub¬ 
bles and make sure 
there are no leaks in 
the system. Keep some 
liquid in the funnel 
throughout the experi¬ 
ment. Fill the gradu¬ 
ated cylinder with water 
and invert it over the 
delivery tube in a 
beaker of water. 

Pour 10 cc. of con¬ 
centrated hydrochloric 
acid into the water in the funnel and allow it to flow into the 
tube until the magnesium has dissolved. When this has oc¬ 
curred, drive all the gas from the system into the graduated 
cylinder by pouring water into the funnel and allowing it to 
flow through the tubes. 

Bring the gas in the cylinder to atmospheric pressure (?) and 
read the volume.-A Calculate the volume the gas would occupy 
at standard conditions. Should you make a correction for 
vapor pressure of water? From the weight of a liter of hydro¬ 
gen at standard conditions calculate the weight of hydrogen 
you obtained. Arrange your data in a table as in Exercise 11. 

From your data calculate the weight of magnesium required 
to displace a gram atom of hydrogen 4i What is this weight of 
magnesium called? Compare this with the theoretical weight 
(equation) and calculate your percentage error. \ 

What is the equivalent weight of aluminum? Of oxygen? 

Of sulfuric acid? Compare the amounts of hydrogen displaced 
by gram atoms of sodium, magnesium, and aluminum. What 
is the relation between the equivalent weight and the atomic 
weight of an element? 

f > /*< 

























t 


26 


GENERAL CHEMISTRY 


EXERCISE 16 


•% 


WATER 


if 


BOILING AND FREEZING POINTS — 
PURIFICATION Ah -,)V- ; ' 

. (L 


Textbook: 108-109; 114-118 


(a) Effect of Solutes. — Place 20 cc. of water in a 100-cc. 
beaker or flask and heat this until it boils. Regulate the flame 
to keep it boiling gently. Immerse the bulb of a thermometer 
in the boiling water and read the temperature finally reached (?). 
Remove the thermometer and saturate the boiling water with 
common salt (about 8 g. needed). When this solution is boiling 
steadily read the maximum temperature reached. How does 
the use of salt affect the time required to cook eggs, vegetables, 
etc., in water? What external condition affects the boiling 
point of a liquid? What is the vapor pressure of water at 
100° C.? What is the effect of a dissolved substance upon the 
vapor pressure of the solvent? 

In the same manner NaCl and other soluble substances lower 
the freezing point of water. If any crushed ice or snow is avail¬ 
able this statement should be verified. Can you think of any 

practical application of this 





Fig. 15 


fact? Explain the use of 
“anti-freeze” mixtures in 
automobile radiators in 
winter. 

(i b ) Preparation of Dis¬ 
tilled Water. — Fit a small 
flask with a one-hole stopper 
and prepare a condensing 
tube as shown in Figure 15. 
The condenser may be made 
in spiral form if a bit of 
skill at glass working is 
possessed. To 40 cc. of 
water add ^ 'cc. of copper 
sulfate solution and •erdtke 


amount of sodium chloride solution. Mix well and test three 
2 -cc. portions with NH 4 OH (?), with BaCl 2 solution (?), and 
with AgN0 3 solution (?). These tests show respectively the 

















Os} p**' &***»***L^ 

JULJ&s . 

LABORATORY EXERCISES 27 

presence of copper, a sulfate, and a chloride. Put the remainder 
of the solution in the flask. Discard the first 5 cc. of the dis¬ 


tillate, and collect 10 cc. more. Make^ie ^ho|hree # |est£^on^ 
small portions of the distillate (?)* $ fykj£X> V " V 

T) Can all substances be removed by distillation? Explain your 
answer. * 71 ^ tc ¥ fLap**# 3 7 

4-4 t W* fan*** ■ $ 0 L / x- !3*r€ L *~ ! J 

EXERCISE 17 


SOME PROPERTIES OF WATER — SOLUTIONS 


Textbook: 119; l^O-l^l 

| 

(a) Combination with Anhydrides. — Moisten a small piece 
of quicklime (calcium oxide) with water and test with litmus 
paper (?). Equation. Recall from Exercise 9 the union of 
water with other basic and acidic oxides. 

( b ) As a Medium of Chemical Action. — Pulverize a gram of 
lead nitrate in a mortar and heat the powder in a porcelain dish 
to dry it. Thoroughly clean the mortar and pestle and pul¬ 
verize a gram of potassium iodide. Heat this powder to dryness 
in another dish or crucible. When both substances are cool mix 
them in one dish (?). Now add a drop of water to the mixture 
of dry salts (?). Add more water and write the equatio n 
for the reaction. Why is water necessary for this chemical 


action?, 

(c) Temperature Effect of Solution. — Draw a beaker of 
water from the tap and measure it^ temperature with a ther¬ 
mometer. In a test tube place 2 g. of calcium chloride and dis¬ 
solve it by shaking with the least amount of water. Note the 
temperature of the calcium chloride solution (?). In the same 
manner dissolve 2 gV of ammonium chloride and take the tem¬ 
perature of the solution (?). Compare the temperatures of the 
solutions with that of the water used. This temperature effect 
is due to the Heat of Solution which may be positive or nega¬ 
tive for different substances. 


4+JufKJ'- 

% i* 


n i. „ L u A 







GENERAL CHEMISTRY 

EXERCISE 18 

HYDROGEN PEROXIDE 


Textbook: 121-125 


(a) Preparation. — Pour into a small beaker 90 cc. of water 
and 10 cc. of dilute sulfuric acid. To this add, in small portions, 
abouDSkg. *of sodium peroxide, stirring vigorously. The solution 
should now be acid (test?), if not, add more sulfuric acid 
diluted as above. Write the equation. Recdh-fche -action of 
Wflrter alone.on sodium peroxide — Exercise 6 (c). 

( b ) Test. — To 5 cc. of this hydrogen peroxide solution add 
about 3 cc. of ether, 2 or 3 drops of potassium dichromate solu- 


overthomiTr^ shake it. The color ' 

the ether layer is used as a test for hydrogen peroxide or, con¬ 
versely, for a dichromate. 

(c) Stability. — To 10 cc. of the solution add a small amount 
of manganese dioxide (?). Test the evolved gas with q glowing 
splinter (?). Is hydrogen peroxide a stable compound? Write 
the equation for its decomposition. What conditions accelerate 
its decomposition? Why is hydrogen peroxide usually sold in 
brown bottles? 


(d) As a Reducing Agent. — Dilute a few drops of potassium 
permanganate solution with 5 cc. of water. Add this solution 
in small portions with shaking to 10 cc. of hydrogen peroxide 
solution. Whak_§as^is.-ovolved? Refer to the textbook for the 
equatiop. What is the role of potassium permanganate in this 
reaction? 

(e) As an Oxidizing Agent. — Dilute 1 cc. of lead nitrate 
solution to 10 cc. and add some hydrogen sulfide solution. 
Shake the tube to suspend the black precipitate (?), pour some 
of this into another tube and add hydrogen peroxide. What 
happens to the black precipitate? Write the equations. 

(/) Evaluation of a Solution of Hydrogen Peroxide. — The 
strength of a given solution of hydrogen peroxide may be de¬ 
termined by two methods: 

(1) The volume of oxygen gas which can be obtained from a 
definite volume of the hydrogen peroxide solution may be 
measured. For example, the 3 per cent commercial product is 



jq^K y ^ 71 ° 3 

LABORATORY EXERCISES 29 

commonly referred to as a 10-volume solution because one 
volume of this solution will liberate ten volumes of oxygen gas. 

If the oxygen evolved in part (c) of tills exercise had been 
measured, this method would have been illustrated. 

(2) If a measured quantity of a solution of potassium per¬ 
manganate of known strength is added to a definite volume of 
hydrogen peroxide solution until a pink color just persists in 
the solution, the strength of the hydrogen peroxide may be 
calculated from the quantity of the permanganate solution used. 

The solution of potassium permanganate on the shelf contains 
15.8 g. of the salt per liter of solution. Dilute 20 cc. of this 
solution to 100 cc. and calculate the strength of this diluted 
solution. Obtain a burette (see Fig. 20) and fill the burette to 
the zero mark with the permanganate solution you have pre¬ 
pared, first making sure the air bubbles have been displaced 
from the rubber connection and tip by bending the tip upward 
and allowing a little of the solution to flow. 

Pour exactly 10 cc. of your hydrogen peroxide solution into a 
beaker, add about 3 cc. of dilute sulfuric acid diluted with 
about 20 cc. of distilled water, and carefully add some of the 
permanganate solution from the burette. Stir the solution in 
the beaker constantly and run the permanganate solution into 
this until a faint pink color persists after stirring. The last 
drop of permanganate solution added should cause this change 
in color. Record the volume of permanganate solution used 
and check this titration with another 10-cc. portion of the hy¬ 
drogen peroxide solution. 

From your data calculate the weight of potassium perman¬ 
ganate which was reduced by 10 cc. of the hydrogen peroxide 
solution, and from this obtain the weight of hydrogen peroxide 
in 10 cc. of your solution and in a liter of such solution. What 
is the percentage of hydrogen peroxide in the solution? What 
weight of hydrogen peroxide should have been produced from 
2 g. of sodium peroxide? If this quantity had been produced, 
what would be the percentage of hydrogen peroxide in your 
solution? 


30 


GENERAL CHEMISTRY 


EXERCISE 19 

THE LAW OF MULTIPLE PROPORTIONS 

Textbook: 125-126 

Quantitative 

A. Refer to Exercise 11 and repeat exactly the directions for 
weighing and heating the materials, except substitute potassium 
perchlorate for potassium chlorate and do not collect the oxygen 
which is evolved. 

Arrange your data from this experiment exactly as you did 
the weight data in Exercise 11. Calculate from the two sets 
of data the weight of oxygen which is combined with one gram 
of potassium chloride and with one gram molecule of potassium 
chloride in potassium chlorate and in potassium perchlorate. 

Show from your results that a simple, whole-number ratio 
exists between the weights of oxygen combined with a fixed 
weight of potassium chloride in the two compounds. 

State the law exemplified by this experiment. Recall or 
obtain from your textbook the formulas for five oxides of nitro¬ 
gen and show clearly, by means of atomic weights, how these 
five compounds illustrate the law just stated. 

B. * In a dry hard-glass test tube place about one gram of 
cupric bromide and weigh the tube and contents accurately. 
Clamp the tube in a position slightly inclined to the horizontal 
and fit it with a one-hole stopper (either cork or rubber) bearing 
a glass tube which dips into a flask containing a little water. 
The end of the tube should not dip below the surface of the 
water, but should end about one inch above it. Heat the cupric 
bromide in the tube, gently at first, then more strongly, until 
bromine fumes are no longer evolved. During this operation it 
is necessary to warm the upper portion of the tube somewhat, 
in order to prevent the condensation of bromine. When the 
tube has cooled, remove the stopper. Tilt the tube so that any 
bromine fumes which remain in it may escape, and weigh the 
tube and its contents. Now pour 2 cc. of concentrated nitric 
acid into the test tube, replace the stopper and delivery tube, 

♦This experiment was contributed by Dr. John C. Bailar, Jr., of the University of 
Illinois. Cf. Jour. Chem. Educ. 6, 1760 (1929). 


LABORATORY EXERCISES 


31 


and boil the liquid gently until all of the acid has evaporated, 
and the residue is dry and black. Replace the one-holed stopper 
by one having two holes; through one hole should pass a tube 
reaching a little more than half way to the bottom of the test 
tube, through the other a tube which extends not more than 
half an inch beyond the inner surface of the stopper. Connect 
one of these to a hydrogen generator; when a test shows that 
all air has been expelled from the apparatus heat the material 
in the test tube until it is completely reduced to metallic copper. 
Allow the tube to cool, while the stream of hydrogen is still 
passing through it, and weigh it again. The data may be re¬ 
corded in the following manner: 

1. Weight of test tube and cupric bromide. 

2. Weight of test tube and cuprous bromide. 

3. Weight of test tube and copper. 

4. Weight of bromine in the cupric bromide. 

5. Weight of bromine in the cuprous bromide. 

6. Ratio of the two amounts of bromine (4 divided by 5). 

State the chemical change which occurs at each step in this 
experiment, writing the equation if possible. Explain clearly 
how your results exemplify the Law of Multiple Proportions. 


EXERCISE 20 

SOLUBILITY — SUPERSATURATED SOLUTION 

Textbook: 135-139 

(a) Influence of Temperature. — To 5 cc. of water in a test 
tube add'd!! cc. of lead nitrate solution and mix by shaking. 
Prepare a similar dilute solution of sodium chloride and add 
half of each solution to a third test tube (?). Heat the remain¬ 
ing halves of the solutions to boiling and mix them (?).. Cool 
this mixture by immersing the test tube in^beakenof y at^ jff)y. 
Heat the other tube containing the preci pixave (:). Wnat* is 
the precipitate? Equation. What can you say of the solu¬ 
bility of this substance? Recall the relation of temperature to. 
solubility of NaCl and KNO3 studied ind£x©reise-4.^ * 

(. b ) Influence of Size of Particles. — Using two small squares 
of paper of exactly equal size counterbalance on the pans of a 


r S ^ 

4 


~ ** fi . 










32 


GENERAL CHEMISTRY 


chemical balance a single crystal (not too large) of copper sul¬ 
fate or potassium dichromate with some of the same substance 
previously powdered in a mortar. Place the two samples in 
test tubes, add equal amounts of water (10 or 15 cc.), shake the 
two tubes equally and measure the time required for each 
sample to completely dissolve. State the factors which influence 
rate of solution. 

(c) Supersaturation. — Weigh out 5 g. of Glauber’s salt 
(Na 2 SO 4 -10H 2 O) and place it in a test tube with 4 cc. of water. 
Warm and shake the tube until all the salt is dissolved. Set 
the test tube in a beaker of cold water or cool it under the tap. 
If no crystals have appeared when the solution is cool, you have 
a supersaturated solution. Now drop into the solution a crystal 
of the salt (?). What kind of solution do you now have? How 
could you determine whether a solution is unsaturated, satu¬ 
rated or supersaturated at a given temperature? 

EXERCISE 21 

HYDRATES 

Textbook: 120; 139-1J+0 

(a) Test for Water of Hydration. — Heat a gram of each of 

the following crystals in dry test tubes and notice whether an 
appreciable amount of water is given off in each case: sodium 
chloride; copper sul fate: sodium phosphate; potassium di¬ 
chromate. If any are found to be hydrate ymte their complete 
formulas. "h 7? 

( b ) Efflorescence and Deliquescence. — Place a few granules 
of calcium chloride on a watch glass or glass plate, some clear 
crystals of Glauber’s salt on another, and a clear crystal of 
copper sulfate on another. Set these away in your desk until 
the next laboratory period, then note and explain any changes. 
Define a deliquescent and efflorescent substance and illustrate 
from your experiment. 

Put a gram of granulated calcium chloride in the bottom of a 
test tube, place a thin tuft of cotton over this, and on the cotton 
lay a clear crystal of copper sulfate. Close the tube tightly 
with a stopper and set away until the next period. Compare 
the appearance of the copper sulfate with that of the sample 
left on a watch glass. Explain. 


LABORATORY EXERCISES 


33 


(c) Hydration. — Mix 10 g. of plaster of Paris ((CaS0 4 )2 * H 2 0) 
and 10 cc. of water in a small beaker, stirring the mixture until 
a thin paste is formed. Pour this into a filter paper fitted in a 
funnel and allow the plaster to “set.” After an hour or more 
examine the product. It is gypsum (formula?); write the equa¬ 
tion. (See Textbook: p. 610.) 


EXERCISE 22 

FORMULA OF A HYDRATE 
Quantitative 

Heat a clean porcelain crucible to redness on a clay triangle, 
supported by the ring stand. Allow the crucible to cool and 
weigh it accurately. Put into it about one gram of gypsum and 
weigh it again. Heat the crucible and gypsum as before to red¬ 
ness for 5 minutes, allow it to cool and weigh the crucible and 
contents again. From the amount of gypsum used and the loss 
in weight, calculate the per cent water of hydration in the 
original compound. Tabulate your data. 

Gypsum is a hydrate of calcium sulfate. Using the atomic 
weights of the elements involved and your own data, calculate 
the correct formula of the hydrate. The following proportion 
is suggested for the solution of this problem: 

Molecular weight of CaS0 4 _ Weight of CaS0 4 left 
x times mol. wt. of H 2 0 Loss in weight. 

Solve for x to obtain the number of molecules of water in a 
molecule of the hydrate. The result should be very nearly a 
whole number. An y variation is due to impurity of the original 
material or to errors in your work. Now write the complete 
formula of gypsum and calculate the theoretical per cent of 
water of hydration. Compare this with your experimental 
result and determine your percentage error. 

Other hydrates, such as borax and crystallized barium 
chloride, may be substituted for gypsum in this experiment. 
CuS0 4 • 5H 2 0 is likely to give incorrect results due to partial 
decomposition of the sulfate into CuO and SO3 if the tempera¬ 
ture becomes too high. 




32 


GENERAL CHEMISTRY 


chemical balance a single crystal (not too large) of copper sul¬ 
fate or potassium dichromate with some of the same substance 
previously powdered in a mortar. Place the two samples in 
test tubes, add equal amounts of water (10 or 15 cc.), shake the 
two tubes equally and measure the time required for each 
sample to completely dissolve. State the factors which influence 
rate of solution. 

(c) Supersaturation. — Weigh out 5 g. of Glauber’s salt 
(Na 2 SO 4 *10H 2 O) and place it in a test tube with 4 cc. of water. 
Warm and shake the tube until all the salt is dissolved. Set 
the test tube in a beaker of cold water or cool it under the tap. 
If no crystals have appeared when the solution is cool, you have 
a supersaturated solution. Now drop into the solution a crystal 
of the salt (?). What kind of solution do you now have? How 
could you determine whether a solution is unsaturated, satu¬ 
rated or supersaturated at a given temperature? 


EXERCISE 21 


HYDRATES 

Textbook: 120; 129-140 

. $M r " 

(а) Test for Water of Hydration. — Heat a gram of each of 

the following crystals in dry test tubes and notice whether an 
appreciable amount of water is given off in each case: sodium 
chloride; copper sul fate: sodium phosphate; potassium di¬ 
chromate. If any are found to be hydratpswrite their complete 
formulas. 3 > * /f ' t 

(б) Efflorescence and Deliquescence. — Place a few granules 
of calcium chloride on a watch glass or glass plate, some clear 
crystals of Glauber’s salt on another, and a clear crystal of 
copper sulfate on another. Set these away in your desk until 
the next laboratory period, then note and explain any changes. 
Define a deliquescent and efflorescent substance and illustrate 
from your experiment. 

Put a gram of granulated calcium chloride in the bottom of a 
test tube, place a thin tuft of cotton over this, and on the cotton 
lay a clear crystal of copper sulfate. Close the tube tightly 
with a stopper and set away until the next period. Compare 
the appearance of the copper sulfate with that of the sample 
left on a watch glass. Explain. 


LABORATORY EXERCISES 


33 


(c) Hydration. — Mix 10 g. of plaster of Paris ((CaS0 4 )2 • H 2 0) 
and 10 cc. of water in a small beaker, stirring the mixture until 
a thin paste is formed. Pour this into a filter paper fitted in a 
funnel and allow the plaster to “set.” After an hour or more 
examine the product. It is gypsum (formula?); write the equa¬ 
tion. (See Textbook: p. 610.) 


EXERCISE 22 

FORMULA OF A HYDRATE 
Quantitative 

Heat a clean porcelain crucible to redness on a clay triangle, 
supported by the ring stand. Allow the crucible to cool and 
weigh it accurately. Put into it about one gram of gypsum and 
weigh it again. Heat the crucible and gypsum as before to red¬ 
ness for 5 minutes, allow it to cool and weigh the crucible and 
contents again. From the amount of gypsum used and the loss 
in weight, calculate the per cent water of hydration in the 
original compound. Tabulate your data. 

Gypsum is a hydrate of calcium sulfate. Using the atomic 
weights of the elements involved and your own data, calculate 
the correct formula of the hydrate. The following proportion 
is suggested for the solution of this problem: 

Molecular weight of CaS0 4 _ Weight of CaS0 4 left 
x times mol. wt. of H 2 O Loss in weight. 

Solve for x to obtain the number of molecules of water in a 
molecule of the hydrate. The result should be very nearly a 
whole number. Any variation is due to impurity of the original 
material or to errors in your work. Now write the complete 
formula of gypsum and calculate the theoretical per cent of 
water of hydration. Compare this with your experimental 
result and determine your percentage error. 

Other hydrates, such as borax and crystallized barium 
chloride, may be substituted for gypsum in this experiment. 
CuS0 4 • 5H 2 0 is likely to give incorrect results due to partial 
decomposition of the sulfate into CuO and SO3 if the tempera¬ 
ture becomes too high. 




34 


GENERAL CHEMISTRY 


EXERCISE 23 

COLLOIDAL SOLUTIONS 

Textbook: 147-155 

Colloidal dispersions (“ solutions ”) are systems in which the 
dispersed (“dissolved”) phase consists of particles considerably 
larger, in general, than the molecules in a true solution and 
somewhat smaller than the particles in a suspension which can 
be separated by filtration or settling. It is obvious, therefore, 
that colloidal dispersions may be prepared by: (1) Dispersion 
methods which involve breaking up larger aggregates into par¬ 
ticles of colloidal dimensions and (2) Condensation methods by 
which substances in true solution are caused to unite and form 
colloidal particles. 

Colloidal particles do not ordinarily settle out, are not re¬ 
moved by filtration, are not visible in a microscope, but their 
presence may be detected by an ultramicroscope due to the 
Tyndall effect. Colloidal particles are electrically charged and 
may be coagulated by the addition of an electrolyte. That ion 
of the electrolyte which bears the charge opposite to that of the 
colloidal particles causes this coagulation, and the effectiveness 
of such an ion increases with its charge or valence. 

(a) Sulfur. — Place a few milligrams of powdered sulfur in 
a test tube, add about 2 cc. of alcohol and shake the tube at 
intervals for five minutes to dissolve the sulfur. It is not neces¬ 
sary that all the sulfur be dissolved. Now pour the alcoholic 
solution of sulfur into about 15 cc. of water in a test tube (?). 
Keep all of the colloidal solutions which you prepare, exhibit 
them to your instructor at the end of this exercise and preserve 
them in corked test tubes for at least several days in order to 
observe their stabilities. 

Prepare a sample of colloidal sulfur by the interaction of a 
dilute solution of sodium thiosulfate with very dilute sulfuric 
acid. If the sulfur obtained by this method tends to settle out, 
the solutions should be diluted still further before mixing. 

Colloidal sulfur may also be produced by oxidizing a solution 
of hydrogen sulfide by simply exposing it to the air or by add¬ 
ing to it a small quantity of acidified potassium permanganate 
solution or other oxidizing agent. 


LABORATORY EXERCISES 


35 


( b ) Silver. — Add ammonium hydroxide solution to a small 
volume of dilute silver nitrate solution until the precipitate 
which first forms is redissolved (?). Now add to this a small 
quantity of a reducing agent such as glucose solution, alcohol, 
or formaldehyde. Warm this preparation until a faint but defi¬ 
nite color develops (?). 

(c) Silver Iodide. — Add 7 cc. of 0.1 normal solution of silver 
nitrate to 20 cc. of 0.05 normal potassium iodide solution (?). 
Consult your instructor with regard to the preparation of these 
solutions from the solutions on the shelf if the latter have a 
higher concentration. 

Repeat this preparation after first adding 1 cc. of a one per 
cent solution of gelatin to each portion of the reactant solutions. 

Prepare a saturated solution of sodium chloride by shaking 
10 g. of the salt with 25 cc. of water in a small flask until no 
more of the salt will dissolve. Determine the minimum volume 
of this solution which must be added to 10 cc. of your first 
preparation of silver iodide in order to produce an evident 
effect of coagulation (?). Add this quantity of the sodium 
chloride solution to 10 cc. of the silver iodide “solution” con¬ 
taining gelatin (?). Explain the protective effect of the gelatin. 
Mention several other protective colloids which might have 
been used. State an example of the use of a protective colloid 
in a commercial product. 

(i d ) Hydrated Iron Oxide. — Heat 50 cc. of water to the 
boiling point; while it is gently boiling add 5 drops of a solution 
of ferric chloride and continue to boil the solution for several 
minutes (?). Add 5 drops of the ferric chloride solution to 
50 cc. of cold water and compare the intensity of color in the 
two preparations. What reaction has occurred in the hot 
solution? If the cold solution is permitted to stand, the color 
will gradually deepen somewhat, showing that the same reac¬ 
tion is occurring in this solution but more slowly. Explain how 
dialysis would aid in carrying this reaction toward completion 
in the cold solution. 


36 


GENERAL CHEMISTRY 


EXERCISE 24 

GELS 

Textbook: 153-15 ^ 

(a) Soap Gels. — Place one gram of soap in each of two 
small beakers. To one of these add 19 cc. of water and to the 
other add 49 cc. of water. Warm each preparation (do not 
boil) until the soap is uniformly dispersed. Permit both solu¬ 
tions to cool to room temperature and let them stand undis¬ 
turbed for an hour or longer. Do either or both of the solutions 
gel? What is the percentage concentration of soap in each 
solution? Is a dispersion of soap in water a true solution or a 
colloidal solution? 

(i b ) Silica Gels. — Dilute 25 cc. of water glass (egg-preserva¬ 
tive grade may be used) with 75 cc. of water and mix the so¬ 
lution thoroughly. To 10 cc. of this solution add a drop of 
phenolphthalein solution and determine approximately the 
quantity of dilute acetic acid (about normal concentration 
should be used) required to decolorize this solution. Add a 
slightly greater quantity of the acetic acid, all at once, to 
another 10-cc. portion of the water-glass solution, mix quickly 
and let this preparation stand until it gels. If the gel forms 
too quickly so that the solutions can not be mixed uniformly, 
the solutions should be diluted still further before mixing. By 
a few trials the proper concentrations to produce a clear, uni¬ 
form gel can be determined. 

What is the composition of water glass (Textbook: p. 426)? 
What reaction occurs when acetic acid is added to water glass? 
Equation. 

Heat a piece of the silica gel in an evaporating dish until no 
further change occurs (?). What is the final product? 

Warm the soap gel prepared in (a) but do not evaporate it 
to dryness (?). Gels may be classified with respect to the 
effect of heat as reversible and non-reversible. Upon this basis 
classify the gels you have prepared. Mention several other 
gel-forming substances in addition to those used in this exercise. 

(c) “Silicate Garden.” — Place the remainder of the water- 
glass solution prepared in ( b ) in a small beaker — at least 50 cc. 
of the solution should be used for this experiment — and drop 




/ 


LABORATORY EXERCISES 37 

into this solution several small crystals of any of the following 
salts: cobalt nitrate, ferric chloride, nickel nitrate, copper 
nitrate or chloride, zinc nitrate or chloride. Permit this prep¬ 
aration to stand for at least an hour. Observe and explain, 
in terms of osmotic pressure, the growths obtained in your 
“silicate garden” (Textbook: p. 145). 

How does a gel differ from a gelatinous precipitate such as is 
obtained when ammonium hydroxide is added to a solution of 
an aluminum or ferric salt? 



%\%- 






ii' 


-L pH m 


EXERCISE 25 


DOUBLE DECOMPOSITION OR METATHESIS 

Textbook: HI 

When two compounds react in solution without any valence 
changes occurring, two other compounds may be completely 
produced if: 

(a) One of the products is insoluble and precipitates. 

(i b ) One of the products is a gas and escapes. 

(c) One of the products is -sligfetty ionize4— e-.g., H 2 0. 

Mix the following solutions in pairs, using about 5 cc. of the first 
solution and adding the second until no further change occurs: 

1. Sodium chloride and silver nitrate. 

2. Barium chloride and dilute sulfuric acid 

3. Sodium carbonate and acetic acid. 

4. Copper nitrate and sodium hydroxide. 

5. Ferric chloride and ammonium hydroxide. 

6. Precipitate from (4) and hydrochloric acid. 

7. Precipitate from (5) and nitric acid. 

Write the equations and indicate in each case the condition 
(a, b or c above) which caused the reaction to go to completion. 

Complete the following type equations: 

An acid plus a base gives — 

An acid plus a salt gives — 

A base plus a salt gives — 

A salt plus a salt gives — 

Illustrate each type with examples other than those given in the 
above experiment. 





38 


GENERAL CHEMISTRY 


EXERCISE 26 

PREPARATION AND PROPERTIES OF CHLORINE 

Textbook: 174.-191 

(a) Preparation. — Set up a flask as in Figure 16 with a de¬ 
livery tube extending down into a dry bottle. Cover the bottle 
with a piece of cardboard or heavy paper having a small hole 
through which the delivery tube passes. (The remainder of 
the apparatus shown in Figure 16 is not used in this exercise.) 

Calculate the weight of manganese dioxide needed to prepare 
2 1. of chlorine by the oxidation of hydrochloric acid. Place 
this weight of manganese dioxide in the flask and add 15 cc. of 
concentrated hydrochloric acid through the thistle tube. Be 
sure the end of the thistle tube dips into the acid. 

Collect and cover with glass plates three bottles of chlorine, 
heating the generator if necessary, and then allow the chlorine 
to bubble into 50 cc. of water in a small flask to prepare “chlo¬ 
rine water.” When this becomes colored fill the generator with 
water to stop the reaction. 

( b ) Bleaching. — In one bottle place a piece of wet, colored 
calico; a dry piece of the same material; wet pieces of paper 
bearing print, pencil and ink writing. Be careful to keep the 
dry cloth out of contact with the wet pieces. After a time 
observe and explain what has happened. Indicate by means 
of an equation the role of water in bleaching with chlorine. 

(c) Reaction with Antimony. — Sift a little powdered an¬ 
timony into a bottle of chlorine (?). Equation. Add water and 
note the formation of a white precipitate (SbOCl). Equation. 

(d ) Reaction with Phosphorus. — Ignite a small piece of 
phosphorus as in Exercise 9 (c) and lower it into a bottle of 
chlorine (?). Note whether liquid phosphorus trichloride or 
solid phosphorus pentachloride is formed and write the proper 
equation. Add some water to the product and test the solution 
with litmus paper. Explain, giving the equation. What ele¬ 
ment does chlorine most resemble with regard to its activity in 
combining with many other elements? 

(e) Chlorine Water. — Test some of the chlorine solution 
with blue litmus, noting and explaining the two changes in color 
which occur. Leave a small piece of dark sponge in a tube of 


LABORATORY EXERCISES 


39 


chlorine water for some time (?). Try some chlorine water on 
the end of a glass rod as an “ink eradicator” (?). Would you 
recommend chlorine water to remove ink stains from dyed fab¬ 
rics? What happens when chlorine water is exposed to sun¬ 
light? 

EXERCISE 27 


HYDROGEN CHLORIDE — HYDROCHLORIC ACID 


Textbook: 193-207 


(a) Preparation — Physical Properties. — To a gram of com¬ 
mon salt in a dry test tube add 0.5 cc. of concentrated sulfuric 
acid. Taking care not to breathe in the gas, note its^odor and 

vP color. Blow your breath across the mouth of the tube (?). 

Test the gas with moist litmus paper (?). Bring a rod moistened 
with ammonium hydroxide into the tube (?). Invert the tube 
in a dish of water and explain the rise of water in the tube. 

Write the equation for the reaction and name the products. 7 
What is the relation of this gas to hydrochloric acid in bottles 
on the desk? 

(b) Test for Chlorides. — Add silver nitrate solution to a few 
drops of each of the following, placed in separate test tubes: 
dilute hydrochloric acid; sodium chloride solution; barium 
chloride solution. Describe the results and write the equations. 
Remove the product from one of the tubes to a filter paper and 
expose it for some time to the brightest light available (?). To 
the product in another of the tubes add ammonium hydroxide 
and slihftfe (?), then acidify it with nitric acid (?). State how 
you could now test for chlorides in a water supply. 

(c) Reaction with Metals. — Recall the action of hydrochlo¬ 
ric acid on metals — Exercise — and write one equation 

to exemplify this. 

(i d ) Reaction with Hydroxides.—To 1 cc. of dilute hydro¬ 
chloric acid add 5 cc. of water, mix and divide the solution in 
three test tubes. Add 2 drops of methyl orange to one, 2 drops 
of litmus to the second, and 2 drops of phenolphthalein to the 
third. Record the colors. Carefully add ammonium hydroxide 
to the first^ nqixing well, until the color changes (?); likewise 
add limewatef to the second (?), and sodium hydroxide to the 
third (?). Tabulate the colors of the three indicators in acid 


k 


40 


GENERAL CHEMISTRY 


/ 


and basic solutions. Write the reaction that occurred in each 
tube. What is this type of reaction called? ' 

Add just enough dilute hydrochloric acid to make the solu¬ 
tion containing phenolphthalein colorless and evaporate soipp 
of it to dryness in a porcelain dish. Taste the product (?).; 

(e) Reaction with Oxides. — Place about 0.24^ g. of zinc 
oxide, in a test tube and add some water. Is it soluble? Add 
dilute hydrochloric acid and shake (?). Equation". 

Repeat the experiment using quicklime. Evaporate tile solu¬ 
tion to dryness. What is the residue? Leave it exposed to the 
air of the laboratory for some time. What property does it 
show? 

Place a few milligrams of ferric oxide in a test tube with 2 or 
3 cc. of concentrated hydrochloric acid and warm until solu¬ 
tion occurs. Pour half of this solution into another test tube 
and dilute it with about 20 cc. of water. Drop an iron nail into 
the other half of the solution, allow the reaction to proceed for 
some time, then dilute with about 20 cc. of water. Compare 
the colors of the two solutions, explain the difference and write 
the equations*,. Also write the equation for the action of hydro¬ 
chloric acid upon iron. ? & 

Add some dilute hydrochloric acid to a few milligrams of 
copper oxide in a test tube and warm until solution occurs. 
Equation. Dilute half of this solution and note a gradual 
change in the tone of the color. Compare this with solutions 
of other copper salts on the shelf (?). Drop an iron nail into 
the other portion of the solution. Why does the reaction grad¬ 
ually decrease? Wash and examine the nail (?). Equation. 

(/) Reaction with Oxidizing Agents. — In separate test tubes 
place 0.1 g. of the following: manganese dioxide, potassium 
permanganate, potassium chlorate, potassium dichromate. 
What property do these substances have in common? To each 
add 1 cc. of concentrated hydrochloric acid; note the color and 
odor (Care) of the evolved gas and test it with moist litmus 
paper and with starch-iodide paper. To prepare starch-iodide 
paper thoroughly moisten a pinch of starch with cold water, 
add 50 cc. of boiling water and 1 cc. of potassium iodide solu¬ 
tion, stir well and dip strips of filter paper into this mixture. 
Chlorine liberates iodine which forms a blue compound with 
starch; an excess of chlorine subsequently destroys this color. 
Hence this may be used as an indirect test for chlorine gas. 


LABORATORY EXERCISES 


41 


Summarize the chemical properties of hydrochloric acid as 
follows: 


HCi Reacting with: 

Products in General: 

Type of Reaction: 

Metals 



Basic Oxides 



Hydroxides 

(Bases) 



Oxidizing Agent 




EXERCISE 28 

POTASSIUM HYPOCHLORITE 

Textbook: 209-214 

Calculate the volume of chlorine needed to react with 5 g. of 
potassium hydroxide, then calculate the weight of potassium 
permanganate and the 
volume of concentrated 
hydrochloric acid (spe¬ 
cific gravity, 1.2, con¬ 
taining 40 per cent 
HC1) required to pro¬ 
duce this volume of 
chlorine. 

Arrange apparatus as 
illustrated in Figure 16. 

In the flask A place 
approximately one-half 
more than the calcu¬ 
lated weight of potas¬ 
sium permanganate. 

Use a small bottle for B and add to it about 10 cc. of water. 
Have the tube from the chlorine generator dip into this; its 
purpose is to remove hydrogen chloride (very soluble) from the 
gas. Dissolve 5 g. of potassium hydroxide in 15 cc. of water, 
dilute 5 cc. of this solution to 20 cc. and reserve it in a test tube, 





































42 


GENERAL CHEMISTRY 


place the remainder of the concentrated solution in the small 
beaker C. 

Pour one-half more than the calculated quantity of concen¬ 
trated hydrochloric acid through the thistle tube and allow the 
chlorine produced to pass into the concentrated solution of 
potassium hydroxide until the solution no longer reacts alkaline 
to litmus paper but bleaches it. Note that the solution becomes 
warm. Write equations for the two reactions that have oc¬ 
curred in the solution. Set aside to crystallize and reserve the 
crystals for the next exercise. 

Allow chlorine to bubble into the dilute solution of potassium 
hydroxide but do not saturate it (litmus test?). Equation. 
Acidify some of this solution and test its bleaching power with 
litmus paper, ink and colored cloth (?). 

State the difference in the products obtained when chlorine 
is passed into dilute and into concentrated solutions of potas¬ 
sium hydroxide. 


EXERCISE 29 

POTASSIUM CHLORATE 

Textbook: 216-218 

Filter off the crystals prepared in the previous exercise and 
wash them twice with 2-cc. portions of water. Test some of the 
filtrate with silver nitrate solution (?). What product does this 
indicate? Dry the crystals between pieces of filter paper and 
try the following tests: 1. Dissolve a few crystals in water, 
acidify the solution with dilute HC1 and compare its bleaching 
power with that of the hypochlorite solution (?). 2. Treat a 
few crystals with concentrated hydrochloric acid (?). 3. Heat 
the remainder of the crystals with a bit of manganese dioxide 
in a small test tube and test for oxygen (?). 

EXERCISE 30 

POTASSIUM PERCHLORATE 

Textbook: 219-220 

Heat 5 g. of potassium chlorate in a test tube until it melts 
and regulate the heat to keep it just above its melting point for 


LABORATORY EXERCISES 


43 


at least 15 minutes. At this temperature you may note a small 
amount of decomposition to form oxygen (equation?), but the 
principal reaction is the formation of the perchlorate (equation?). 

When the tube is cool remove the solid, grind it in a mortar 
and transfer it to a beaker. Dissolve this mixture of salts by 
adding portions of water, heating it to boiling with stirring and 
pouring off the solution into another beaker. Repeat this 
process until the whole mass has been brought into solution. 
Allow the solution to cool and note the formation of crystals. 
What are they? What other salts are left in solution? What 
must be the relative solubility of the salts? 

Carefully decant the mother liquor, bring the crystals on to 
a filter paper and wash them a few times with 2-cc. portions of 
water. Dry the crystals between filter papers; place a small 
quantity of them in a test tube; place a similar quantity of 
potassium chlorate in another tube and add concentrated hydro¬ 
chloric acid to both (?). Which oxy-acid of chlorine is the most 
stable? Will an acidified solution of potassium perchlorate 
bleach litmus paper or cloth? Will potassium perchlorate give 
oxygen when heated? From the relative stabilities of the two 
salts would you expect the chlorate or the perchlorate to de¬ 
compose at the lower temperature? 


EXERCISE 31 

BLEACHING POWDER 

Textbook: 21^-215 

Place a small lump of quicklime in a 100-cc. beaker and care¬ 
fully slake it by adding just enough water to cause the quick¬ 
lime to crumble. Note the development of heat. Equation. 
Spread the slaked lime over the bottom of the beaker, generate 
some chlorine in a test tube (materials?) fitted with a stopper 
and right-angled delivery tube and allow the chlorine to pass 
into the beaker over the slaked lime. Note that absorption of 
the gas occurs and write the equation. The product CaCl(OCl), 
a “mixed salt,” is the salt of what two acids? How would you 
name it to indicate this fact? 

To a small portion of the bleaching powder add dilute sulfuric 
acid (?). Equation. 


44 


GENERAL CHEMISTRY 


Mix some bleaching powder with 50 cc. of water, acidify this 
with a few drops of sulfuric acid and dip a piece of colored calico 
into the solution until a change is noted. Also try the action of 
this solution on news print and on a spot of fountain-pen ink on 
cloth (?). What is the actual bleaching agent in this solution? 
By means of an equation show the purpose of adding the sulfuric 
acid. 


EXERCISE 32 

IONIZATION 

Textbook: 232-236 

As indicated in Figure 17 allow the carbon rods attached to 
an electrical circuit to dip into a U-tube containing the following 
solutions separately: 

(a) Distilled water 

(b) Dilute hydrochloric acid 

(c) Sodium hydroxide solution 

(d) Sugar solution 

(e) Salt solution 

(/) Acetic acid solution 

(i g ) Alcohol solution 

( h ) Ammonium hydroxide solution. 

The U-tube and carbon electrodes must be rinsed with water 
after each trial before introducing the next solution. 

If the electric light in the 
circuit glows when the 
switch is closed, conduct¬ 
ance is shown, and the 
brightness of the light in¬ 
dicates the relative conduc¬ 
tivity of the solution. What 1 
classes of substances con¬ 
duct electricity in solution? 
Why? Indicate the ions, if 
any, that are present in 
each case. Define an elec¬ 
trolyte. Which substances 
in the above solutions are electrolytes? 


+ 










LABORATORY EXERCISES 


45 


EXERCISE 33 

IONIZATION AND CHEMICAL ACTION 

Textbook: 232-238 

(a) H-Ion Concentration and Rate of Reaction. — Prepare a 
series of mixtures of concentrated sulfuric acid and water in six 


bes as follows: 



Tube Number 

Cc. Sulfuric Acid 

Cc. Water 

1 

10 

0 

2 

8 

2 

3 

6 

4 

4 

4 

6 

5 

2 

8 

6 

0 

10 


In mixing the acid and water always pour the acid into the water 
and mix carefully. When the mixtures have cooled to room 
temperature, drop a piece of zinc into each of the six tubes. 
Observe carefully and compare the rate of hydrogen evolution 
in the several tubes. The reaction by which zinc liberates 
hydrogen may be expressed as 

Zn + 2H + —» Zn ++ + H 2 

and consists in the transfer of two electrons from each atom of 
zinc to two hydrogen ions. Since the concentration of zinc is 
constant, the rate of reaction is naturally governed by the con¬ 
centration of hydrogen ions. What can you conclude from 
your results as to the relative concentrations of hydrogen ions 
in the various mixtures? Are there any hydrogen ions in pure 
water? 

Heat the first tube, containing zinc and concentrated sulfuric 
acid, and note that a gas is evolved (odor?) which is not hydro¬ 
gen. It is due to the reaction of zinc with molecules of sulfuric 
acid. Equation. 

(6) Effect of the Solvent. — Recall or test the action of a 
water solution of hydrogen chloride upon zinc and upon cal¬ 
cium carbonate (marble). Equations. 

Obtain or prepare 10 cc. of a dry solution of hydrogen chloride 


48 


GENERAL CHEMISTRY 


(d) Test for Bromide. — From parts (a) and (6) of this exer¬ 
cise deduce several characteristic tests for a bromide. Add 
silver nitrate solution to the solution of a bromide and test the 
solubility of the precipitate in ammonium hydroxide (?). Com¬ 
pare this with the chloride test — Exercise 27 ( b ). 


EXERCISE 37 

IODINE 

Textbook: 251^0 

(a) Preparation. — Repeat the first paragraph of Exercise 
36 (a) using potassium iodide solution in place of the bromide (?). 

Grind together 1 g. of potassium iodide and 2 g. of manganese 
dioxide. Transfer the mixture to a porcelain dish and moisten 
it with a solution of 1 cc. concentrated sulfuric acid in 2 cc. of 
water. Invert a clean dry funnel over this dish and gently heat 
it on a wire gauze (Fig. 18). Moist pieces of filter paper may 
be placed on the outside of the funnel to cool the surface upon 
-which the iodine will condense. Note the color of the vapor 
and of the W$fta¥s. Write the equation for the reaction. When 
a considerable amount of iodine has 
been collected scrape it into a clean 
watch glass and use it for the experi¬ 
ments below. 

(6) Physical Properties. — Heat a 
crystal of iodine in a test tube. Does it 
melt? Sublime ? Define sublimation. 

Place a few iodine crystals in each 
three test tubes. Add 10 cc./m Water to 
the fir^t and shake well, 1 cc. of alcohol 
to the second, and 1 cc. of solution of 
potassium iodide to the third (?). Is 
iodine as soluble in water as are chlorine 
Fig. 18 ^,1 anc j bromine? What is a solution of 

iodine in alcohol called? 

(c) Starch Test. — Prepare a starch suspension by boiling a 
pinch of starch in 10 cc. of water. Test the water solution of 
iodine with some of this (?). To some starch suspension add 
1 cc. of a solution of potassium iodide (?); now add chlorine 

















LABORATORY EXERCISES 


49 


water (?). Is starch a test for iodine in a compound? Dilute 
some of the starch suspension, add a drop of the alcoholic solu¬ 
tion of iodine and dilute further until a transp^reMblue is 
obtained. Heat the contents of the tub^xtr bcmnfgthe 
color should reappear on cooling. How may this test be used 
conversely to test for starch in foods, baking powder, face 
powder, etc.? 

(d) Hydriodic Acid. — To a few milligrams of potassium 
iodide in a dry test tube add several drops of concentrated 
sulfuric acid. Warm gently if necessary and note the color and 
odor of the gaseous products. Equations. What property of 
hydrogen iodide is shown here? Compare Exercise 36 (&). 
How may pure HI be prepared? 

0) Test for an Iodide. — From the foregoing experiments 
state, how you would test for an iodide. To solutions of a 
-v chloride, a bromide, and an iodide add a solution of mercuric 
nitrate and from your results state how you could distinguish 
the three halides. 

EXERCISE 38 

HYDROGEN FLUORIDE (HYDROFLUORIC ACID) 


Textbook: 260-262 

Cover one side of a glass plate with a thin layer of paraffin 
by warming the plate uniformly at some distance above a flame 
and rubbing it with a piece of solid paraffin. When the paraffin 
has cooled, draw a figure or some letters with a pointed instru¬ 
ment in the center of the waxed surface, exposing the glass. 

Place 5 g. of calcium fluoride (fluorspar) in a lead dish (store¬ 
room), moistenJhis with 2 or 3 cc. of concentrated sulfuric acid, 
stirring the mixture with a match stick, and cover the dish with 
the glass plate — waxed surface downward. ( Caution: Do not 
breathe the fumes as they are very poisonous.) 

Set this preparation away in the hood until the following day 
or later, then remove the paraffin and examine the plate for the 
etched design. The paraffin may be removed by warming the 
plate and wiping it with a towel or paper. Construct the equa¬ 
tions for the preparation of hydrogen fluoride and for the etch¬ 
ing of glass. Glass is a mixture of silicates^ CaSi0 3 may be 
taken as representative. 

Of what materials must containers for hydrofluoric acid be 




I 


50 


GENERAL CHEMISTRY 


made? Can hydrofluoric acid be oxidized to fluorine and water? 
What is the action of fluorine on water? Equation. 

EXERCISE 39 

PREPARATION OF PURE SODIUM CHLORIDE 
FROM ROCK SALT 

Grind about 35 g. of rock salt in a mortar with 100 cc. of 
water. Transfer the whole to a beaker and warm with stirring 
to complete the solution of the salt. This will make an approxi¬ 
mately saturated solution. 

Dissolve 2 g. of sodium carbonate in 2 or 3 cc. of water and 
add this to the hot salt solution to precipitate the calcium and 
magnesium (?). Equation. Stir, allow the precipitate to settle, 
add a few drops of sodium carbonate solution and notice whether 

or not more precipitate forms. 
If it does, add more sodium 
carbonate solution to remove 
completely the calcium and mag¬ 
nesium. 

Filter the hot solution, allow 
the filtrate to cool, and pass in 
HC1 gas until a considerable 
amount of salt has precipitated. 
The HC1 gas is conveniently 
supplied by heating some concentrated hydrochloric acid in a 
test tube provided with a one-hole stopper and delivery tube, 
bent to pass downwards. An inverted funnel should be at¬ 
tached to the delivery tube and so adjusted that the mouth of 
the funnel dips just below the surface of the salt^olution (Fig. 19). 

The effect of HC1 in precipitating the salt from a saturated 
solution may be explained by the following equilibrium: 

NaCl (solid) <=±NaCl (in solution) Na + + Cl - . Passing HC1 
gas into the solution is equivalent to adding Cl - ions, hence the 
equilibrium is driven toward the left (see Textbook: 236-238). 

When a sufficient amount of salt has precipitated, allow it to 
settle and decant the clear liquid. Using successive 10 cc. por¬ 
tions of dilute hydrochloric acid, wash the salt by decantation 
until the washings no longer give the barium chloride test for 
sulfates (?). 

Transfer the solid to a filter and allow it to drain for some 













LABORATORY EXERCISES 


51 


time. Dry the salt by heating it carefully in a clean porcelain 
dish, stirring it with a glass rod. The salt should be perfectly 
white, finely crystalline, and should have a pleasant taste. 
Since the chlorides of calcium and magnesium have been re¬ 
moved, it is not hygroscopic and will remain dry in a damp 
atmosphere. Submit the product to your instructor for ap¬ 
proval. 


EXERCISE 40 
ATOMIC STRUCTURE 

Textbook: 279-299 

Write clear definitions for the terms electron , proton , atomic 
number , isotopes. Contrast the proton and the electron with 
respect to charge, size and mass (weight). How many protons 
and how many electrons are there in a hydrogen atom? in a 
hydrogen molecule? in an oxygen atom? in a molecule of water? 
What are the atomic or molecular weights of these units of 
matter? Is the mass of each atom or molecule due principally 
to the protons or to the electrons? Does it appear that the 
total number of protons in an atom or molecule is numerically 
the same as the atomic or molecular weight? If this is true, 
what can you say as to the total number of electrons in an atom 
or molecule of weight n? 

Draw a diagram consisting of concentric circles to represent 
the structure of each of the following atoms: 



Atomic Number 

Atomic Weight 

(a) Oxygen. 

. 8 

16 

(6) Sulfur. 

. 16 

32 

(c) Calcium. 

. 20 

40 

(d) Selenium. 

. 34 

76 

(e) Selenium. 

. 34 

80 

(/) Krypton. 

. 36 

84 


Let the inmost circle represent the nucleus of the atom. Indi¬ 
cate the number of protons and electrons contained in this, 
using appropriate symbols for proton and electron. Let each 
succeeding circle represent a shell of electrons and show the 
number of electrons which it contains. Show clearly the rela¬ 
tion between atomic number and the composition of the nucleus, 
and between atomic number and the number of electrons out¬ 
side the nucleus. 








4 


52 


GENERAL CHEMISTRY 


Define positive and negative valence in terms of electrons. 
From your atomic diagrams derive the positive or negative 
valence of each element you have pictured. Upon this basis 
predict the formula of the compound which selenium will form 
with calcium and the formula of a compound of selenium with 
oxygen. Explain the mechanism by which each combination 
will take place. 

Describe the similarity which is apparent in your structures 
for oxygen, sulfur and selenium. What is the significance of 
this with regard to the relative positions of these elements in 
the periodic system? 

Is the difference between two isotopes of an element due to 
a difference in the nuclei or to a difference in the electron shells? 
How do you account for the fact that the atomic weight of 
chlorine is approximately 35.5? Zinc has four isotopes having 
masses of 64, 66, 68 and 70. How many different zinc chlorides 
having the formula ZnCl 2 can be formed and what will be the 
molecular weights of the different species? 


EXERCISE 41 


THE PER CENT OF OXYGEN IN THE AIR 

Textbook: 307-308 

Determine the capacity of the graduated cylinder by filling 
it level full with water and withdrawing the water by means of 
a pipette to the 100-cc. mark. The last por¬ 
tion withdrawn can be measured in the 10-cc. 
graduate. 

Obtain a piece of white phosphorus in a dish 
of water, push the end of a piece of copper or 
iron wire through the phosphorus, place this 
upright over a beaker of water, and immedi¬ 
ately cover it with the graduated cylinder. 
Figure 20 shows the arrangement of the appa¬ 
ratus. Be sure the open end of the graduate 
extends considerably below the surface of the 
water in the beaker and clamp the cylinder in 
— this position. Using a form such as that sug¬ 
gested below, make a record in your notebook 
of the beginning conditions of the experiment and then set the 


















LABORATORY EXERCISES 


53 


apparatus away in your desk until the next laboratory class. 
At that time adjust the levels of water inside the cylinder and 
in the beaker and take the readings necessary to complete your 
record. 

Compare your result with the accepted value for the per cent 
by volume of oxygen in the air and calculate your percentage 
error. Write the equations to show what has happened to the 
oxygen in the cylinder. 


Table of Data 

Volume of air. cc. 

Temperature. °C. 

Barometric pressure. mm. 

Vapor pressure of water. mm. 

Corrected volume of air (S. C.)... cc. 

Volume of gases left. cc. 

Temperature. °C. 

Barometric pressure. mm. 

Vapor pressure. mm. 

Corrected volume (S. C.). cc. 

Volume of oxygen removed (S. C.) cc. 

Per cent oxygen in air. % 


EXERCISE 42 


AMMONIA — PREPARATION; PROPERTIES 


Textbook: 326-337 




(a) Preparation. — In separate test tubes place 0.25 g. of (1) 
ammonium chloride, (2) ammonium nitrate, and (3) ammonium 
sulfate. To each tube add 2 cc. oEsodium hydroxide solution. 
Warm each mixture, notice the mfe apd test the issuing gas 
with moist htimi^feper and with a piece of filter paper moist¬ 
ened with concentrated hydrochloric acid (?). Grind a little 
quicklime or slaked lime with a small amount of an ammonium 
salt and test for ammonia alftbove (?) ^Equations. What 


general statement can you make concerning the preparation of 
ammonia? How would you test for an ammonium salt? 

(, b ) Solubility; Aqua Ammonia. — Test for the presence of 
ammonia gas above a solution of ammonium hydroxide (?). 
What is your conclusion as to the stability of this compound? 












GENERAL CHEMISTRY 




54 


Fit a small flask with a stopper bearing a tube extending 
through in both directions as in Figure 21. Pour 2 cc. of con¬ 
centrated aqua ammonia into the flask and warm 

® this gently for a moment. Have a drop of water in 
that end of the tube which will extend into the 
flask, quickly insert the tube and stopper and invert 
the flask above a beaker of water to which a few 
drops of phenolphthalein have been added. Ex¬ 
plain all that occurs. 

(c) Combination with Acids. — In one test tube 
place 0.25 g. of sodium chloride and add a few drops 
of concentrated sulfuric acid. In another test tube 
place 0.25 g. of ammonium chloride mixed with an 
equal weight of powdered lime; heat this mixture 
and bring the two test tubes mouth-to-mouth (?). 
Write the equations. 

Write equations to show the combination of am¬ 
monia with five acids and name the salt formed in each case. 

(, d ) Dissociation of Ammonium Salts; Diffusion. — Place 
0.5 g. ammonium chloride in a dry test tube and clamp it in an 
inclined position. In the mouth of the tube place moist pieces 
of red and blue litmus. Gently heat the ammonium chloride, 
note its apparent sublimation and observe the color changes of 
the litmus. Compare the relative weights of the two gases 
formed from ammonium chloride and explain the order of the 
litmus changes. 


Fig. 21 


EXERCISE 43 


NITRIC ACID — PREPARATION AND PROPERTIES 

Textbook: 34-1-350 

(a) Preparation — Oxidizing Action. — Place a gram of po¬ 
tassium nitrate in a test tube, add 1 cc. of concentrated sulfuric 
acid and insert a small plug of woolen yarn in the mouth of the 

! in a holder or clamp and warm 

with 


tube. Support the test tube 


mixture until the vapor of nitric acid comes in contact 




the 

the wool (?). 

What other acids may be prepared by the' Action of sulfuric 
acid on their salts? Why may nitric acid be prepared by this 
method? 










LABORATORY EXERCISES 


55 m 




(b) Action on Metals. — Place in segarate test^|ubes a small 
piece of the following metals: zfeSf co^per. ffn, n§T-(nail), 
magnesiu m and aluminum. Add to each enough dilute nitric 
acid to cover the metal and note the relative speed with which 
each metal is acted upon — or the order in which the tubes 
begin to show the evolution of a gas. Add a few drops of 
concentrated nitric acid to any tube that fails to show any ^ 
action. 

According to your results arrange the metals in a column 
showing the relative ease with which they are oxidized by nitric 
acid. For comparison arrange these same metals in another 
parallel column according to their positions in the Electromo¬ 
tive Series (back of book or chart in classrooms). Note any 
difference of position in the two arrangements (?). Note: 
Nitric acid acts upon aluminum to form a continuous coating 
of insoluble A1 2 0 3 which prevents any further action. Does 
this suggest a method of removing stains from aluminum 
utensils without harm to the metal? 

h I^^ ener al> what are the products of the action of nitric acid 
on metals? ^Contrast the acid in this respect with sulfuric (di¬ 
lute and concentrated) and witJ^hydjochloric acid, f 

(c) Decomposition. — HeatT^rop? of concentrated nitric 
acid in a dry test tube and note the decomposition products. 
Equation. In the same maprji^. a ip crj^ls pfepppp 
mtrate or lead nitrate. Indicate how tnis method might be 
used to prepare the pure oxides of these metals. For what 
nitrogen compound does this suggest a method of preparation? 

(d) Test. — An interesting test for nitric acid or nitrates, 
referred to as the “brown ring test,” may be carried out as 
follows :> Dissolve one or two crystals of potassium nitrate in 
3 cc. of water, add an equal volume of ferrous sulfate solution 
and mix well. Slightly incline the test tube and cautiously 
pour concentrated sulfuric acid down the inside of the tube. 

The acid sinks to the botft>m and a colored layer forms where 
the two liquids meet. 

Outli ne jin alternative test fc^-a nitrate, using concentrated 
sulfurie^cm ; alA copper. 

(e) Summary of Properties. — Tabulate the properties of 
nitric acid as you did those of hydrochloric acid in Exercise 27. 
What is the essential difference in the chemical properties of 
these two acids? 






56 


GENERAL CHEMISTRY 




EXERCISE 44 



REDUCTION PRODUCTS OF NITRIC ACID 

Textbook: 339-34-0; 346-348 

ty) Nitric Oxide. — Place a gram of copper turnings in a test 
tube fitted with a stopper and bent delivery tube and arrange 
to collect several test tubes of gas by displacement of water. 
Add 3 cc. of dilute nitric acid (sp. gr. 1.2), clamp the tube in 
position and, if necessary, warm it to start the action. Explain 
why the gas in the generating tube is brown at first but becomes 
colorless as it passes into the collecting tube. 

Thrust a glowing splinter into a test tube full of this gas (?). 
Compare the per cent of oxygen in this gas with the per cent 
of oxygen in air and explain the result just obtained. 

Place another test tube full of nitric oxide mouth-to-mouth 
with a test tube full of air (?). Can you note any indications 
of temperature or volume changes in the reaction that occurs? 
Add 3 or 4 cc. of water to one of the tubes, place your thumb 
over the mouth of the tube and shake it until the gas becomes 
colorless. Remove the thumb for a moment and repeat the 
procedure (?). Test the solution with litmus paper and write 
equations to explain the changes observed. 

Calculate the volume of nitric oxide which may be obtained 
by the action of nitric acid on a gram of copper. 

(6) Nitrogen Dioxide. — Put a few drops of concentrated 
nitric acid in a test tube and drop into this a small piece of 
copper (?). Compare this with the preparation of nitric oxide 
(?). What other method of preparing nitrogen dioxide is sug¬ 
gested by reactions in (a)? Indicate how this gas could also be 
prepared from the nitrate of a heavy metal. 

(c) Ammonia. — Place a few very small pieces of zinc in a 
100-cc. beaker, add 2 cc. of water and 2 cc. of dilute nitric acid. 
Allow the action to continue for|5 minutes, then make the solu¬ 
tion alkaline with sodium jjjjBSfoxide solution and cover the 
beaker with a watch glass having a piece of wet red litmus paper 
adhering to its lower surface. Warm the mixture slightly if 
necessary, note the effect on the litmus and the odor of the gas 
in the beaker (?). Why did the ammonia not escape as it was 
formed? Write the equations. 

■ m 

m 



LABORATORY EXERCISES 


57 $ 


Upon what two factors does the nature of the reduction 
product of nitric acid depend? Indicate six possible reduction 
products of nitric acid, each containing nitrogen with a differ¬ 
ent valence. Write separate equations to show the reduction 
of nitric acid by means of hydrogen to each of these products. 


EXERCISE 45 


NITROUS OXIDE 

Textbook: 338 

Heat 2 g. of ammonium nitrate in a clean test tube provided 
with a delivery tube as in Exercise 44 (a). The water over which 
the gas is collected should preferably be warm. (Why?) The 
ammonium nitrate should be heated gently until it melts and 
then caused to decompose slowly by holding the Bunsen burner 
in the hand and waving the flame under the test tube to pro¬ 
duce a steady evolution of gas. If this substance is heated too 
rapidly it may explode. 

Collect several test tubes full of the gas by displacement of 
water. The first should be discarded as it will contain air. 
Using a test tube full ofthe 
solubility in cold water (7). 

end of a wood splinter into a test tube full of the gas (?). Now 
introduce a brightly glowing spark into another tube full of the 
gas (?). How dqes nitrous oxid e compare with oxygen as a 
supporter of combustion? What products would be formed if 
phosphorus were burned in this gas? What is the common 
name of nitrous oxide? - 

Compare the percentage of oxygen in the three oxides of 
nitrogen with which you have experimented. Does the one 
containing most oxygen support combustion best? y. Explain. 
Calculate the weight of oxygen in a liter of nitroUs oxide and 
in a liter of nitric oxide, y ) 1/ * 

Contrast the decomposition products of ammonium nitrate 


,£oS ? device a method of testing its 
'Introduce a faint spark on the 






0 


i 

with the products obtained by heating ammonium nitrite. ^ 
Indicate the valence of each nitrogen atom in ammonium 
nitrate and in ammonium nitrite and show how these valences 
change when each compound is decomposed. 

$ v 


a 


to 





58 


GENERAL CHEMISTRY 


EXERCISE 46 

NITRITES — NITROUS ANHYDRIDE 

Textbook: 340-341 

Heat 0.5 g. of sodium nitrate or potassium nitrate in a test 
tube until it decomposes, testing the decomposition product 
with a glowing splinter (?). Compare this result with the de¬ 
composition of two other nitrates — Exercises 43 (c) and 45. 
After the residue in the test tube has cooled add 1 cc. of dilute 
sulfuric acid (?). State how this would distinguish a nitrite 
from a nitrate — both will respond to the “brown ring test.” 

Melt a gram of sodium nitrate in a crucible, add a gram of 
“test” lead and continue heating the mixture for some time, 
stirring it with a rod or file. What compound of the metal is 
formed? What happens to the nitrate? When the crucible is 
cool add a small amount of water and dissolve the salt by warm¬ 
ing and stirring. Pour this solution into a test tube, cool it and 
add dilute sulfuric acid. Note the color of the solution and of 
the gases evolved when the solution is warmed. 

What is the action of nitrous anhydride on cold water? What 
happens when this solution is warmed? Equations. 


EXERCISE 47 

SULFUR — ALLOTROPIC FORMS 

Textbook: 352-358 

(a) Changes on Heating. — Fill a test tube one-fourth full of 
sulfur and heat it slowly until it boils, carefully noting all the 
intermediate changes. Allow the tube to cool, noting the re¬ 
versal of the changes and save it for (c). 

(b) Rhombic. — Shake 0.5 g. of sulfur in a test tube with 3 
or 4 cc. of carbon disulfide and filter the solution into a clean 
watch glass. {Caution: Avoid flames.) Allow the solution 
to evaporate; watch the formation of crystals, noting their 
color and shape, and indicate the shape of a good crystal in 
your notes by means of a drawing. 

(c) Monoclinic. — Slowly reheat the tube of sulfur from (a) 




LABORATORY EXERCISES 


59 


until the mobile stage of molten sulfur is just reached. Pour a 
few drops of this into a beaker of water and a small amount on 
a cold plate or dish. Compare the crystals formed with those 
prepared in ( b ). 

(d) Amorphous (Plastic). — Continue heating the molten 
sulfur until the dark viscous stage is reached or even to the 
boiling point and then pour it into a beaker of cold water (?). 
Compare this product with the drops chilled in (c). Is the 
plastic form of sulfur stable? Leave it in your desk for a few 
days and examine it again (?). Test its solubility now and 
later in carbon disulfide (?). Above what temperature must 
molten sulfur be heated in order to form plastic sulfur when 
chilled? Why? 


EXERCISE 48 

SULFUR — COMBINATION WITH METALS; 

HYDROGEN SULFIDE 

Textbook: 357-359 

(а) With Iron and with Copper. — Prepare two small closed 
tubes about 12 cm. long as in Exercise 6 (6). Place a small 
amount of sulfur in the bottom of each tube and about 2 cm. 
above the sulfur place a wad of copper thread or fine wire in 
one tube and some iron picture wire in the other. In each case 
heat the tube about the metal until it is red-hot and then boil 
the sulfur. Notice any indications of chemical action as the 
sulfur vapor comes in contact with the hot metal (?). 

Allow the tubes to cool, remove and examine the products 
and treat each sulfide with dilute sulfuric acid. If a gas is 
evolved notice its odor (?). Equations. 

(б) With Silver — Removal of Tarnish. — Clean a small 
silver coin by washing it in hot sodium hydroxide solution 
followed by rinsing. Place a few drops of ammonium polysul¬ 
fide, (NH 4 ) 2 S x , on the coin and note the formation of silver sul¬ 
fide. The tarnishing of silver in the household is principally 
due to the formation of this substance from sulfur compounds 
in eggs, perspiration, and illuminating or cooking gas. The 
tarnish may be removed without harm to the silver by placing 
it in a hot salt solution in contact with a piece of aluminum 
e.g., an aluminum pan. 


60 


GENERAL CHEMISTRY 


Pour 10 cc. of water into a beaker and add enough sodium 
chloride to make a saturated solution. Place the tarnished 
coin in this solution in contact with a piece of aluminum and 
boil the solution. If the method does not work the aluminum 
must be removed and cleaned by abrasion or by boiling it in a 
solution of sodium hydroxide for a moment. (Cf. Textbook: 
585.) 

EXERCISE 49 

HYDROGEN SULFIDE — PROPERTIES AND USES 

Textbook: 358-362 

(a) Preparation. — Place a small piece of ferrous sulfide in a 
test tube and add 1 cc. of dilute hydrochloric or sulfuric acid (?). 
Cautiously note the odor of the gas and apply a flame to the 
mouth of the tube (?). Write equations to show the production 
and combustion of the gas. 

(i b ) Precipitation of Sulfides. — Put 1 cc. of the solutions of 
metal salts fisted below into separate test tubes and add 5 cc. 
of hydrogen sulfide solution to each (?). If no precipitate forms 
in certain tubes add ammonium hydroxide to those tubes (?). 
Finally test the solubility of each precipitate in dilute hydro¬ 
chloric acid and tabulate your results as indicated below: 


Solution of: 

Precipitate with 
H 2 S? Color? 

Ppt. with H 2 S 4* 
NH 4 OH? Color? 

Soluble in HCl? 

Pb(N0 3 ) 2 



-AyO 

Mn(N0 3 ) 2 




CuS0 4 



y t • 

ZnS0 4 




N a 2 S0 4 


9 // 

j iJUalS. 

FeS0 4 



^ ... / 11 

SbCl 3 



_ 

CaCl 2 

k 1 



AsC1 3 
























LABORATORY EXERCISES 


61 


CLASSIFICATION 

Three Groups of Sulfides 

1. Precipitated in acid solution: 

(Insoluble in dilute acid.)-- 

2. Precipitated in alkaline solution: 

(Soluble in dilute acid, insoluble in water.) - 

3. Not precipitated in either acid or alkaline solution: 

(Soluble in water or decomposed by water.) - 

Write equations for any chemical changes which occurred. 

(c) Separation of Two Metals as Sulfides. — To 5 cc. of water 
add 3 or 4 drops of an antimony chloride solution and just 
enough hydrochloric acid to dissolve the white precipitate (?). 
Now add a few drops of a solution of ferrous sulfate or ferrous 
chloride and 5 cc. of hydrogen sulfide water. Shake the mix¬ 
ture well and filter it. Make the filtrate alkaline with ammo¬ 
nium hydroxide and filter through another filter paper. What 
remains on each filter paper? Write the equations. As classi¬ 
fied in (6), into what groups do these sulfides fall? 

(i d ) As a Reducing Agent. — Put 1 cc. of a ferric chloride 
solution in a test tube, t add some hydrogen sulfide water and 
warm gently. Note “the chfihge in the color of the solution and 
the formation of a fine precipitate. Write the equation. 

Burn some sulfur in a bottle of air, add 10 cc. of water, shake 
and pour the solution (?) into a test tube. Add hydrogen sul¬ 
fide water to this and explain the result. What changes in 
valence occur in this reaction? 

To 1 cc. of a solution of potassium dichromate add a small 
quantity of hydrochloric acid and then hydrogen sulfide water 
until the solution changes color (?) and a milky precipitate 
results. Chromium is reduced from the dichromate to CrCl 3 . 
Indicate the valence changes and balance the equation. 

K > ** 9 & 

EXERCISE 50 

SULFUR DIOXIDE — PREPARATION; PROPERTIES 

Textbook: 366-369 

(a) Preparation. — Recall one method already employed in 
making sulfur dioxide (?). Obtain 0.5 g. sodium sulfite or 





*/*>*<ty 

62 


t (Ua** **> ^ a 

GENERAL CHEMISTRY 


sodium bisulfite (formula?) in a test tube and add a few drops 
of dilute sulfuric acid. Test the evolved gas with moist blue 
• litmus and note its odor by cautiously wafting some of it 
toward your nose. 

Place 3 g. of copper turnings and 5 cc. of concentrated sul¬ 
furic acid in a test tube fitted with a one-hole stopper and a 
delivery tube bent to pass downward when the test tube is 
clamped in an inclined position. Gently warm the acid until 
it reacts upon the copper and collect two test tubes full of the 
gas by conducting it into the test tube, held in an upright posi¬ 
tion, until a litmus test or the odor of the gas at the mouth of 
the tube indicates that it is full. Stopper the tubes and save 
them and the generating tube for the following experiments. 

Indicate by means of equations three different methods for 
preparing sulfur dioxide. 

(6) As a Bleaching Agent. — Introduce some flower petals, 
pieces of green leaf or grass into one test tube of sulfur dioxide 
and again stopper the tube. After a time note any change that 
has occurred. Compare the bleaching action of sulfur dioxide 
with that of chlorine in regard to the type of chemical action, 
permanency and power. 

(c) As an Acid Anhydride. — Add 3 or 4 cc. of water to the 
other tube of sulfur dioxide, quickly cover the mouth of the 
tube with the thumb and shake the tube. Note and explain 
the suction felt when the thumb is now removed (?). Test^the 
solution with litmus paper. What has been formed? Note the 
odor of the solution. What is your conclusion with regard to 
the stability of the product? What other anhydride does sulfur 
form? Write equations to show the action of water on the two 
anhydrides. 

Pour 5 cc. of a solution of sodium carbonate into a test tube 
and pass sulfur dioxide into the solution from the generator pre¬ 
pared in (a). After a short time the solution begins to effervesce. 
What is the gas? Equation. Continue passing in sulfur dioxide 
until effervescence ceases and evaporate the resulting solution to 
dryness in a porcelain dish. What remains? Treat this residue 
with a few drops of dilute sulfuric acid (?). 

(, d ) As a Reducing Agent. — Saturate 10 cc. of water in a 
test tube with sulfur dioxide by passing in the gas at such a rate 
that the bubbles can be counted. Note that the bubbles of 
gas decrease in size and often disappear before they reach the 




*7 

■^VVv. 


* * * 


LABORATORY EXERCISES 


63 


^fU 


top of the solution. Explain. Continue passing in sulfur diox¬ 
ide until bubbles of the gas rise through the solution with little 
or no decrease in size (?). 

Divide this solution among three test tubes. To one portion 
add 2 cc. of bromine water (?) and to a second portion add, drop 
by drop, a solution of potassium permanganate until a delicate 
pink color persists in the solution. Now add a few drops of a 
solution of barium chloride and some dilute hydrochloric acid 
to each of the three tubes (?). For what ion is the barium ion 
a test? Write equations to show all that you observed. 




*4 


EXERCISE 51 

SULFURIC ACID 




Textbook: 370-377 


(a) Preparation. — Burn some sulfur, held in a deflagrating 
spoon, in each of two wide-mouthed bottles for a minute or 
o and cover the bottles. 

Pour 10 cc. of water into one of the bottles, shake it and test 
he solution with litmus (?). Add a few drops of a solution of 
irium chloride and some dilute hydrochloric acid. To the 
bottle of sulfur dioxide add a few drops of concentrated 
acid. Place the bottle mouth-to-mouth with a bottle of 


nr 


air and pour the fumes back and forth (?). Then add 10 cc. of 
water, shake and add barium chloride solution and dilute hydro¬ 
chloric acid as before. 

Compare the results in the two bottles and indicate by 
mean%of an equation the action of the nitric acid. In what 
;pects 


rv 


respe^es does this experiment resemble the lead-chamber process 
for the manufacture of sulfuric acid? 

(i b ) Temperature and Volume Changes on Dilution. — Care¬ 
fully measure 40 cc. of water in the graduated cylinder and pour 
it into a small flask. Measure exactly 10 cc. of concentrated 
sulfuric acid in the small graduated cylinder and pour the acid 
into the water in the flask. Note the heat produced and ex¬ 
plain why it is safer to pour concentrated sulfuric acid into 
water rather than water into the acid. Cool the solution, 
measure it in the cylinder and calculate the per cent change in 
volume. 


<uU/ 






64 GENERAL CHEMISTRY 

(c) As a Dehydrating Agent. — Try the effect of a few drops 
of concentrated sulfuric acid upon a pinch of sugar, bits of 
wood, paper and cloth. The formula of cane sugar is C 12 H 22 O 11 ; 
indicate the effect of concentrated sulfuric acid upon it. Take 
warning from these observations and avoid spattering sulfuric 
acid on clothing or skin. What would be the result if dilute 
sulfuric acid were spilled on cloth and allowed to dry? 

( 1 d ) As an Oxidizing Agent. — Recall the action of hot con¬ 
centrated sulfuric acid on copper (see Exercise 50). Test its 
action upon a small iron nail or piece of zinc (?). Compare 
these results with the action of dilute sulfuric acid upon the 
same metals. 

(e) Action on Salts. — Recall the action of concentrated sul¬ 
furic acid on fluorides and chlorides (?). What is its action on 
a nitrate? What property of sulfuric acid makes it valuable for 
the preparation of other acids? Why can it not be used to 
prepare hydrobromic and hydriodic acids from bromides and 
iodides? Write the equation for its action upon KBr. 

1(f) Solubility of Sulfates. — Add a few drops of sulfuric acid 
to solutions of salts of lead, calcium, strontium and barium (?). 
Try the barium chloride test on a solution of a sulfate from the 
side shelf and also on a soluble carbonate and soluble phosphate, 
noting the effect first of the barium chloride, and then of the 
addition of dilute hydrochloric acid to the mixture in each of 
the three tubes. 

EXERCISE 52 

SODIUM THIOSULFATE 

Textbook: 377-379 

Prepare sulfur dioxide as in Exercise 50 and pass it into 5 cc. 
of a 10 per cent solution of sodium hydroxide until the solution 
is saturated. The solution should now give an acid test with 
litmus due to the presence of sodium bisulfite. Equation. Add 
more sodium hydroxide to this solution, a few drops at a time, 
with shaking, until the solution is just slightly alkaline. To this 
add 0.5 g. powdered sulfur, boil gently for at least 10 minutes 
and filter. Equations. 

Test a small portion of the filtrate with dilute sulfuric acid; 
observe and explain a gradual change in the appearance of the 



f \<W 



LABORATORY EXERCISES 65 

solution. Equation. Compare this with the action of an acid 
on a sulfite (?). 

Pour about 2 cc. of water into a test tube, add tgtfthis 2 drops 
of silver nitrate solution and 2 drops of a solution of sodium 
chloride (?). Test the solubility of this precipitate in some of 
the thiosulfate solution. State the application of the observed 
result to photography. 

Stain a piece of cloth with a drop of tincture of iodine or dip 
the cloth in chlorine water and note that the cloth retains the 
chlorine odor. Now dip the cloth in some sodium thiosulfate 
solution and note the effect (?). Suggest the practical applica¬ 
tion of. the result. 


• EXERCISE 53 


N*pRMAL SOLUTIONS — EQUIVALENT WEIGHTS 

Textbook: 879-388 

rite in your notebook answers to the following: 

f I* What is a standard solution? A molar solution? A nor- 
fmal solution? An equivalent weight? (Refer to Exercise 15.) 
dlus rate each of these definitions by means of an example. 

% 2. jkow much is a gram equivalent of phosphoric acid? Of 

Wa? 

3. ^A gram molecule of phosphoric acid contains how many 
gram equivalents? How many liters of normal solution would 
it make? It would make 2 1. of what normality? 

4. What volume of a base of 0.25 normality (N/4) will be 
neutralized by 100 c. of a twice normal (2N) acid? 


Problem. — (a) If a liter of N potassium hydroxide solution 
is treated with hydrochloric acid and evaporated, what weight 
of'salt remains? If 99.4 g. of this salt are obtained in this 
way from a liter of potassium hydroxide, what was the normality 
of the base? 

( b ) If 100 cc. of an unknown acid neutralize 150 cc. of the 
base of the above strength, what is the normality of the un¬ 
known acid? 


66 


GENERAL CHEMISTRY 


EXERCISE 54 

NORMAL SOLUTIONS — ACIDIMETRY 

Textbook: 380-382 

Quantitative — Formal Report Required 

Measure 65 cc. of a 10 per cent solution of sodium hydroxide 
into a clean bottle, add enough water to dilute the solution to 
150 cc. and mix well. This solution 
should now be slightly stronger than 

normal. 

Weigh accurately a clean, dry,30-cc. 
porcelain dish. Measure exactly 10 cc. 
of the sodium hydroxide solution into the 
weighed dish by means of a pipet^.f To 
use the pipette, suck the liquid tbove 
the 10-cc. mark, being careful not to 
draw it into your mouth, quickly place 
index finger over the upper end o^lhe 
pipette and by carefully releasing the 
finger allow the liquid to run down until j 
the meniscus is exactly at the j mark^v 
Then allow this 10 cc. of solution] to run 
from the pipette into the dish, touching 
the tip of the pipette to the inner surSBe 
of the dish to remove the last drop. 

Add 5 cc. of dilute hydrochloric acid to the solution in the 
dish, place the dish on a water bath and evaporate the contents 
to apparent dryness. Now cover the dish with a watch glass 
and heat it gently on a wire gauze until no more moisture con¬ 
denses on the watch glass, then heat fairly hot. When the dish 
is quite cool weigh it again. From the weight of the salt ob¬ 
tained, calculate the weight of sodium hydroxide in 10 cc. of 
your solution and from this the normality of the solution. Is 
it necessary to know the exact amount of hydrochloric acid 
added provided an excess was used? Why? 

With a Report Card, provided by your instructor, and a clean, 
dry, 100-cc. graduated cylinder obtain a sample of “unknown” 
acid. Set up a burette as in Figure 22, wash it with water, 
rinse with a small amount of your sodium hydroxide solution, 















LABORATORY EXERCISES 


67 


and fill the burette with this solution. Make sure that all the 
air bubbles have been displaced from the tip of the burette 
and its rubber connection. This may be accomplished by 
bending the tip upward and allowing a little of the solution to 
flow. Finally, adjust the bottom of the meniscus to the zero 
reading of the burette. 

Wash and dry the 10-cc. pipette or rinse it with some of 
the unknown acid. Accurately measure 10 cc. of the un¬ 
known acid in the pipette, and transfer it to a 100-cc. flask 
and dilute it with water to about 50 cc. Add 2 or 3 drops of 
a solution of methyl orange or phenolphthalein. Recall the 
color changes of these indicators — Exercise 27 (d). Now care¬ 
fully run in the sodium hydroxide solution from the burette 
until the end-point is exactly reached. The last drop of solu¬ 
tion added should cause the indicator to change color. The 
solution must be stirred or shaken after the addition of each 
small portion of sodium hydroxide solution. Repeat this titra¬ 
tion several times with 10-cc. portions of the unknown acid and 
average your results. 

| From the volume of the sodium hydroxide solution required 
hnd its normality, calculate the normality of the unknown 
acid — i.e.j how many cubic centimeters of an exactly normal 
acid are equivalent to one cubic centimeter of the unknown acid. 

: 

EXERCISE 55 

PHOSPHORUS 

Textbook: 384-397 

(a) Allotropy. — From a piece of glass tubing prepare a tube 
about 8 cm. long and closed at one end. Place a few milligrams 
of red phosphorus in this tube and heat it until some of the phos¬ 
phorus has distilled and condensed farther up the tube. When 
the tube is cool, cut it under water at this point and remove 
some of the distilled phosphorus by means of a wire or the for¬ 
ceps. ( Caution: White phosphorus ignites in air and if in con¬ 
tact with flesh produces painful burns which heal very slowly. 
It should not be handled with the fingers.) Dry this on a filter 
paper and leave it exposed to the air until it ignites. Contrast 
the two forms of phosphorus with respect to several properties. 


68 


GENERAL CHEMISTRY 


(b) Three Types of Phosphates. — Place about 0.2 g. of 
disodium phosphate, Na 2 HP0 4 - 12H 2 0, in a porcelain crucible 
and heat it to redness for some time. Equation. Allow the 
product to cool and dissolve it in water. ^ this solution add 
a few drops of silver nitrate solution (?).^ 1 ^y 

Heat about 0.2 g. of microcosmic salt, NaNH 4 HPO 4 -4H 2 0, 
in a crucible as above. What is evolved? Equation. Wh|n 
cool dissolve the product and add silver nitrate solution 7 !^), j 

Test a solution of disodium phosphate with silver nitrate 
Ablution (?). Test the solubility of half of the precipitate in an 
excess of a dilute acid and the other half in ammonium hydrox- 
^" ide (?). Write equations for the reactions of silver nitrate with 
sodium pyrophosphate, sodium metaphosphate, and disodium 
orthophosphate. Compare the colors and forms of any pre¬ 
cipitates so produced. 

(c) Phosphoric Acid. — Slowly heat a minute quantity of red 
phosphorus with 1 cc. of concentrated nitric acid in a porcelain 
evaporating dish until brown fumes are no longer evolved. 
When cool take up the residue with 2 or 3 cc. of water, filter if 
necessary and test the solution with silver nitrate solution. 
According to the color of the precipitate (cf. b) which|pho|| 
phoric acid have you obtained? Equation. 

(d) Calcium “ Superphosphate.” — To 5 g. of bone ash (ter¬ 
tiary calcium phosphate) in a porcelain evaporating dish add 
3 cc. of water and 3 cc. of concentrated sulfuric acid. Warm 
the mixture for 10 minutes stirring occasionally. Add 10 cc. of 
water, stir and filter the solution. Neutralize the solution with 
ammonia and then make it just slightly acid with dilute HN0 3 . 
Test the resulting solution with silver nitrate solution (?). Why 
was it necessary to neutralize the excess acid and also to avoid 


an ammoniacal solution — refer to (6)? 

What is the correct chemical name for the “superphosphate”? 
For what is the substance used? Why is it preferable to bone 
ash for this purpose? 


LABORATORY EXERCISES 


69 


EXERCISE 56 

ARSENIC AND ANTIMONY 

Textbook: 399-412 

{Caution: Arsenic and antimony compounds are very poi¬ 
sonous.) 

(а) Reduction of Arsenic Trioxide. — Prepare a tube as in 
Exercise 55 (a). Place a few milligrams of arsenic trioxide in 
the bottom of the tube, cut a piece of charcoal to fit into the 
tube and place this about 1 cm. above the trioxide in the tube. 
Heat the charcoal red-hot and, while keeping it hot, heat the 
arsenic trioxide to cause its vapor to pass over the charcoal. 
What is the black ring formed above the charcoal? Heat the 
black deposit and note whether it melts or sublimes (?). Heat 
some powdered antimony in a small glass tube and compare its 
behavior with that of arsenic. 

(б) Arsenious Sulfide and Ammonium Sulfarsenite. — Dilute 
1 cc. of arsenic trichloride in a ? test tube with water and add 
some hydrogen sulfide^solution (?). Allow the precipitate to 
settle, pour off most of the supernatant liquid and test the 
solubility of the precipitate in a mixture of ammonium hydrox- 
ide^ancLhydrogen sulfide solution, or in a solution, of ammonium 
polysulfide, (NH 4 ) 2 S X . If thd^tt^r is used, ammonium sulf- 
arsenate is formed. Write the equations. Repeat the experi¬ 
ment using lead nitrate solution in place of arsenic trichloride 
and from your results indicate how the sulfides of lead and 
arsenic could be separated. 

(c) Colloidal Arsenious Sulfide. — Boil a small quantity of 
arsenious oxide with 20 cc. of distilled water and filter the solu¬ 
tion. To the filtrate add some hydrogen sulfide solution. What 
reaction do you expect? Is the product visible? Compare with 
the result in (6) when H 2 S was added to a solution of AsC 1 3 and 
explain the difference. To different portions of the colloidal 
solution just obtained add a few drops of the following solutions: 
sodium chloride, hydrochloric acid, acetic acid, sugar. Explain 
any difference in the results and classify as electrolytes and non¬ 
electrolytes the solutions which were added. 

(d) Displacement. — Put 2 cc. of a solution of antimony 
trichloride in a test tube, add an equal volume of water and 



70 


GENERAL CHEMISTRY 


enough hydrochloric acid to dissolve the precipitate (?). Drop 
a piece of tin into this solution and watch the action for a few 
minutes (?). Drop a piece of tin into a solution of arsenic 
trichloride and compare the deposit formed on the tin in the 
two cases. 

Study this type of reaction with a piece of zinc in a solution 
of lead nitrate and a clean iron nail in a solution of copper sul¬ 
fate. Write equations for all the above changes. What metals 
could be displaced from solution by copper? Make a general 
statement with regard to the displacement of a metal from so¬ 
lution by another metal. 

( e ) Amphoteric Character. — Add some ammonium hydrox¬ 
ide to 1 cc. of antimony trichloride solution. Divide the pre¬ 
cipitate into two parts: to one add hydrochloric acid, to the 
other a solution of sodium hydroxide. Write the equations. Is 
Sb(OH) 3 an acid or a base? What kind of oxide is Sb 2 0 3 ? 
What kind of an element is antimony? 

(/) Antimony Sulfide. — To 1 cc. of a solution of antimony 
trichloride add some water and note the hydrolysis. Add Br few 
drops of hydrochloric acid (?) and then add some hydrogen 
sulfide solution (?). This sulfide is similar to As 2 S 3 in its action 
with ammonium sulfide (cf. 6). Equation. Pour off the super¬ 
natant liquid and dissolve the precipitate by adding a few drops 
of concentrated hydrochloric acid. Show that this is a reversible 
reaction. 

EXERCISE 57 

BISMUTH 

Textbook: ^12-415 

(а) Basic Bismuth Nitrate. — Warm 0.1 g. of metallic bis¬ 
muth in about 1 cc. of concentrated nitric acid until solution 
occurs. Equation. To the resulting solution add water, drop 
by drop, until a precipitate forms (?). What is the pharmaceu¬ 
tical name of this product and for what is it used? 

(б) Reduction of the Basic Oxide. — To 1 cc. of a solution of 
bismuth trichloride add some ammonium hydroxide and filter 
off the precipitate. Prepare some sodium stannite, Na 2 Sn0 2 , by 
adding sodium hydroxide to 1 cc. of stannous chloride until the 
precipitate formed at first is dissolved in the excess sodium 
hydroxide. Add this solution to the precipitate of BiO(OH) on 


LABORATORY EXERCISES 71 

the filter paper (?). The reduction product is metallic bismuth. 
Equation. 

(c) Bismuth Sulfide. — Pour 1 cc. of a solution of bismuth 
chloride into a test tube, add water and if a precipitate forms (?) ' M ' 

dissolve it by adding dilute hydrochloric acid. Now add hydro¬ 
gen sulfide solution and compare the precipitate with the cor¬ 
responding compounds of arsenic and antimony. 

Is bismuth more o^less metallic than antimony? Give sev¬ 
eral reasons for your answer. How would you distinguish a 
solution of SbCl 3 from a solution of BiCl 3 ? 

EXERCISE 58 

SILICON 

Textbook: 417-4-30 

(a) Reduction of Silica — Hydrogen Silicide. — Mix a gram 
of fine sand with the calculated amount of magnesium powder 
required to reduce the silica to silicon and to combine with the 
latter to form magnesium silicide. Place the mixture in a test 
tube and heat it gradually until a reaction occurs (?) and then 
heat it as hot as possible for a few minutes. Be careful not to 
point the test tube at anyone while heating it. Allow the prod¬ 
uct to cool, break the tube and cautiously throw some of its 
contents into a small beaker containing 5 cc. of concentrated 
hydrochloric acid (?). Write the equations. 

(b ) Soluble Glass — Silicic Acid. — In a porcelain crucible 
mix one gram of infusorial earth (“kieselguhr,” formula?) with 
2 g. of sodium carbonate, 2 g. of potassium carbonate, and a 
piece of solid potassium hydroxide (about 2 g.). Heat this mix¬ 
ture as hot as you can for at least 15 minutes. When the cruci¬ 
ble is cool, place it with its contents in a small beaker, add 25 cc. 
of water and boil for some time to dissolve the soluble portion. 

Filter this solution and divide the filtrate into two portions. 

To one portion of the filtrate add 5 cc. of concentrated hydro¬ 
chloric acid, drop by drop, stirring the solution (?). Equation. 

To the other portion add 5 cc. of concentrated hydrochloric 
acid all at once (?). Evaporate this solution almost to dryness 
in a dish or beaker (?). If the same product results here as in 
the other portion why did you not see it at first? What is the 
final product that may be obtained by heating this residue? 


72 


GENERAL CHEMISTRY 


What is sodium silicate commonly called? Mention several 
of its uses. 

(c) Glass. — Make a small loop at the end of a platinum 
wire which is fused into a piece of glass tubing at its other end. 
The glass tubing is to serve as a handle for the wire. Heat the 
loop for a moment in the Bunsen flame and dip it into some 
powdered sodium carbonate. Heat again and repeat the opera¬ 
tion until the sodium carbonate forms a clear bead in the loop. 
Touch the bead while hot to some powdered silica and heat the 
bead as hot as possible in order to dissolve the silica. Likewise 
take up some powdered calcium carbonate on the bead and 
continue heating and dipping the bead in the different sub¬ 
stances until a clear fusible bead is obtained. This bead should 
now have a composition similar to that of ordinary glass. Write 
the equations for the reactions which have occurred. 

Grind a few fragments of glass to a powder in the mortar, 
moisten this powder with a few drops of water and add to it a 
drop of phenolphthalein solution (?). Why does not water in 
an ordinary glass vessel give this test? 

Mention three advantages of fused quartz as compared with 
ordinary glass. How may glass or quartz be etched? Write 
equations to illustrate this. (Recall Exercise 38.) 


EXERCISE 59 

BORON 

Textbook: ^31-^3^ 

(a) Boric Acid. — Dissolve 5 g. of borax in 15 cc. of hot 
water in a small beaker. Add the calculated amount of dilute 
sulfuric acid to convert the borax to boric acid, assuming 1 cc. 
of dilute sulfuric acid to contain 0.25 g. H 2 S0 4 . Allow the solu¬ 
tion to cool and examine the product. For what is it used? 
What is the anhydride of boric acid? Of what acid is borax a 
salt? What is its anhydride? Is boric acid more or less hydrated 
than the acid corresponding to borax? 

(b) Borax Glass — Bead Tests. — Make a loop about 3 mm. 
in diameter in the end of a platinum wire as in Exercise 58 (c). 
Heat the loop for a moment in the Bunsen flame and dip it into 
some powdered borax. Heat again until the borax which ad- 


LABORATORY EXERCISES 


73 


hered to the hot wire melts to a clear bead in the loop. Touch 
this bead while hot to some minute particles of cobalt oxide or 
cobalt nitrate and heat again until the particles dissolve in the 
bead. The resulting color is characteristic of cobalt glass. 
Other metallic oxides dissolve in borax glass to give character¬ 
istic colors. The borax bead test may be repeated with com¬ 
pounds of copper, iron, nickel and manganese. In some cases a 
different color is obtained for the same metal in the oxidizing 
and reducing zones of the flame. 

What properties of a “boro-silicate” glass make it superior 
to ordinary glass for chemical apparatus and for cooking ware? 

(c) Tests for Borates. — Place a small quantity of borax or 
boric acid in a porcelain dish, moisten this with some concen¬ 
trated sulfuric acid, and add a few cubic centimeters of alcohol. 
Light the alcohol in the open dish and observe the color of the 
flame as contrasted with the ordinary flame of alcohol. 

Saturate a few cubic centimeters of water with borax and 
add dilute hydrochloric acid to this solution until it is acid to 
litmus. Moisten a piece of turmeric paper with this solution, 
note its color and dry it on a watch glass over a beaker of 
boiling water. When the turmeric paper is dry touch it with 
a drop of sodium hydroxide solution, by means of a glass rod, 
and note the change in color. 


EXERCISE 60 
CARBON 

Textbook: 436-454 

(, a ) Preparation of Charcoal. — Place a small piece of wood 
in an iron crucible having a diameter of about 3 cm. Cover the 
crucible and heat it on a clay triangle for several minutes with 
the full flame of the Bunsen burner. Try to determine whether 
or not the escaping products are combustible. When the cru¬ 
cible is cold examine the residue (?). 

( b ) Preparation of Coke. — Repeat (a) using about 2 g. of 
coal instead of the piece of wood. 

(c) Adsorptive Power of Charcoal. — Dilute a drop of methyl 
violet to 10 cc. (“indelible” pencils contain this dye), add 2 g. 
of animal charcoal (bone black), boil and filter the solution (?). 


74 


GENERAL CHEMISTRY 


Suggest some practical application of this property of finely 
divided carbon. 

(d) Carbon Dioxide — Carbonates. — Fit a test tube with a 
one-hole stopper bearing a delivery tube bent twice at right 
angles so as to pass downward. Place about 3 g. of marble in 
the test tube and add 5 cc. of dilute sulfuric acid. Insert the 
stopper with the delivery tube and allow the gas to bubble into 
2 cc. of sodium hydroxide solution for a short time (?). Now 
add dilute sulfuric acid to this solution (?). Equations. Of 
what acid is carbon dioxide the anhydride? 

Take 5 cc. of clear limewater in a test tube, add an equal 
volume of water and pass carbon dioxide into the solution until 
the turbidity which appears at first has been cleared up (?). 
Boil the solution (?). Explain the relation of this result to the 
formation of a deposit in teakettles and boilers when certain 
hard waters are used. 


EXERCISE 61 


IDENTIFICATION OF A SALT — A REVIEW 


Purpose: 

1. To discover the acid radical present in an unknown 
substance. 

2. To test whether or not the substance is a hydrate. 

3. To test whether or not the substance is an ammonium salt. 

Procedure: 




tv Q i 

k ' v \ 


Devise and write equations for distinguishing reactions 
of salts of the following acids: 

HC1, HBr, HL^HN0 3 , HND 2 , H 2 S H 2 S0 3 , H 2 S0 4 , 

h 2 s 2 o 3 , h 3 po 4 /h 3 bo 3 , h 2 co^^>t 

For example, an^’chloride with concentrated sulfuric acid 
and manganese dioxide gives chlorine gas which may be 
recognized. 

Next devise a test for an ammonium salt and also for a 
hydrate. 

Present this outline to your instructor for approval. Try 
all the identifying tests with known substances from the 
shelf, recording in your notebook such phenomena as 
precipitates (color), gases evolved (color and odor), etc. 


LABORATORY EXERCISES 


75 


(c) Your instructor will now give you a Report Card and you 
may obtain an “unknown” from the storeroom. Ana¬ 
lyze this substance by the tests you have outlined. 

(< d ) Record your results in your notebook and on the Report 
Card and return the latter to the storeroom. 

EXERCISE 62 

A REVIEW OF OXIDATION-REDUCTION 

Equations by the Ion-Electron Method 

Textbook: 201-203 

Define oxidation and reduction in terms of valence change. 
Recall the idea of positive and negative valence arrived at in 
Exercise 40 (cf. Textbook, Chapter 17) and define oxidation 
and reduction in terms of a loss or gain of electrons. 

What is the essential difference between an atom of sodium 
and a sodium ion? Is oxidation or reduction involved when a 
metal ion, such as the cupric ion, is deposited in the electrolysis 
of a cupric chloride solution? Explain. What must happen to 
a sulfur atom in order that it may become a sulfide ion? 

Refer to the third paragraph of Exercise 49 (d) and write the 
valence changes and the balanced equation for the action which 
occurred. Recalling that most acids, bases, and salts are prac¬ 
tically completely ionized, it is apparent that in this reaction 
there are certain ions which are present both before and after 
the chemical change but which have taken no part in the reac¬ 
tion. To which two ions in this reaction does the above state¬ 
ment apply? The equation can, therefore, be formulated from 
the other ions and the electrons which are lost and gained by 
these ions, omitting entirely the unchanged ions. It is most 
convenient to do this in two 'partial equations, one involving 
the loss of electrons, the other the gain of electrons. If this loss 
and gain is equalized by multiplying the partial equations by 
the necessary factors and the equations are added, the electrons 
cancel out and the balanced ionic equation is obtained. 

Illustration: The equation to which we have referred should 
be formulated by this method as follows: 

(R) Cr 2 0 7 = + 14H+ + 6e 2Cr+++ + 7H 2 0 

(O) S= -> S ' + 2e 


76 


GENERAL CHEMISTRY 


Multiplying the oxidation (O) partial equation by 3 and adding 
the result is 

Cr 2 0 7 = + 3S= + 14H+ -> 2Cr+++ + 3S + 7H 2 0. 

This balanced ionic equation has advantages over the usual 
molecular equation in that: 

(1) Only those ions appear which are actually involved in 
the reaction. 

(2) The equation is general — that is, it represents the reac¬ 
tion of any dichromate with any sulfide in the presence of any 
acid. 

By means of the method just outlined write the ionic equa¬ 
tions for the oxidation-reduction reactions which you observed 
in the following experiments: 

Exercise 27 (/) Mn0 2 + Cl - + H + —> Mn ++ 

Mn0 4 " + Cl" + H+ Mn++ + Cl 2 + H 2 0 
C10 3 - + Cl- + H+ -> 

Cr 2 0 7 - + Cl" + H+ -> Cr+++ 

Exercise 28 Cl 2 -f OH - —> Cl - + CIO - + H 2 0 
Cl 2 + OH- -> Cl- + C10 3 - + H 2 0 

Exercise 36 (a) Cl 2 + Br~ —> 

Br- + Mn0 2 + H+ -> 

(b) Br- + H+ + S0 4 = -> 

Exercise 44 (a, b, c) 

Exercise 50 (c) 

Compare the final ionic equation in each case with the mo¬ 
lecular equation which you wrote previously, noting the ions 
which were inactive. 

Write the ionic equation for the action of nitric acid upon a 
sulfide to produce nitric oxide and free sulfur. Translate this 
into a molecular equation for copper sulfide. 

Formulate the ionic equation for the action of a halogen (X) 
with a hot, concentrated base. Translate this into a molecular 
equation for iodine and barium hydroxide. 

Write the ionic equations for the following reactions, stating 
which ions are inactive. 

CaBr 2 + H 2 S0 4 + Na 2 Cr0 4 -> 

FeS0 4 + H 2 S0 4 + HN0 3 -> 

Na 2 S0 3 + HC1 + KMn0 4 -> 


PART II 


THE METALS 
EXERCISE 63 

THE PHYSICAL PROPERTIES OF THE METALS 

Textbook: 522-531 

Refer to the table of physical properties of the metals and 
to your textbook in answering the following questions. 

1. Contrast the physical properties of the metals and non- 
metals in at least three particulars. 

2. State the physical property of the metal concerned which 
makes possible the use indicated: 

gold and silver thread for cloth; gold leaf for lettering; 
aluminum in aeroplane frames; lead for plumbing; 
lead for bullets and shot; copper for trolley wires. 

3. Which metal has the highest melting and boiling points? 
What use is made of the metal because of this fact? 

4. Which metal has a melting (or freezing) point below or¬ 
dinary temperatures? What use of this metal depends 
upon this fact? How does the boiling point of this metal 
compare with the boiling points of other metals? Do you 
know of any appliance in which the vapor of this metal is 
used? 

5. Will the following metals float or sink in mercury: gold, 
silver, platinum, lead, tungsten, iron? 

6. Which is the heaviest metal listed in the table? How does 
its melting point compare with the melting points of other 
common metals? What use for the metal does this suggest? 
Mention other uses of this metal. 

7. How many pounds should a “gold brick’’ weigh if it meas¬ 
ures 20 cm. X 10 cm. X 10 cm. and is real? What would 
it weigh if it were composed of lead with a thin plating of 
gold? (1 lb. = 454 g.; 1 kg. = 2.2 lb.) 

8. Mention two metals which can be melted in a tube placed in 
boiling water. Mention two metals which will float on water. 

77 


78 


GENERAL CHEMISTRY 


9. Which seven metals were known and used in very ancient 
times? Which of the metals commonly used in the house¬ 
hold is of most recent discovery? 

10. Does there appear to be any general relationship between 
the ability of a metal to conduct heat and its ability to 
conduct electricity? Illustrate your answer by means of 
at least two instances. 

11. Which metal is the best conductor of heat and electricity? 
Which metal is most commonly employed for electrical 
conduction? Why do you give a different answer to these 
two questions? 

12. The conductivity comparisons in the table are for wires of 
the same size; if the comparison is made on the basis of 
wires of the same weight for a given length, which metal is 
the best conductor? Suggest some reasons why this metal 
is not more commonly used for electrical conduction. 

13. From the data in the table suggest two or more physical 
reasons why aluminum is a good material for cooking 
utensils. Assuming other conditions the same, will water 
boil more quickly in a vessel of iron, platinum, copper, or 
gold? Give three physical reasons why lead would be an 
undesirable material for a frying pan. 


EXERCISE 64 

PROPERTIES OF ALLOYS 

Textbook: 581-534. 

(a) General. — Alloys are mixtures of two or more metals, 
solid solutions of metals in each other, or, in some cases, definite 
compounds of one metal with another. Alloys are usually pre¬ 
pared by melting together the component metals. The physical 
properties of an alloy may be intermediate to, or greater than, 
or less than the corresponding values of the component metals. 

Referring to the table of physical properties, what do you 
conclude as to the general effect of alloying upon the following 
properties of the metals: density, melting point, conductivity? 
Particular care is taken to remove traces of other metals from 
copper which is to be used for electrical conduction. Can you 
suggest the reason for this? 


LABORATORY EXERCISES 


79 


PHYSICAL PROPERTIES OF THE METALS 


Metal 

Density 
or 3p. Gr. 
(H 2 0 = 1) 

Melting 

Point 

°C. 

Boiling 

Point 

°C. 

Electri¬ 
cal Re¬ 
sistivity 
20°C. 

Thermal 

Conduc¬ 

tivity 

Date 
of Dis¬ 
covery 

Aluminum 

2.7 

658.7 

1800 

2.82 

.504 

1828 

Antimony 

6.62 

630 

1440 

41.7 

.042 

1450 

Arsenic 

5.73 

Sublimes 

450 

33.3 

— 

1694 

Barium 

3.8 

850 

950 


— 

1808 

Bismuth 

9.78 

269.2 

1436 

120. 

.0194 

1450 

Cadmium 

8.65 

320.9 

778 

7.6 

.222 

1817 

Calcium 

1.54 

810 


4.6 


1808 

Chromium 

6.92 

1615 

2200 

2.6 


1797 

Cobalt 

8.72 

1480 


9.7 


1773 

Copper 

8.94 

1083 

2310 

1.72 

.918 

Prehist. 

Gold 

19.32 

1063 

2500 

2.44 

.700 

Prehist. 

Iron 

7.86 

1530 

2450 

10. 

.161 

Prehist. 

Lead 

11.34 

327 

1525 

22. 

.083 

Prehist. 

Magnesium 

1.74 

651 

1120 

4.6 

.376 

1829 

Manganese 

7.42 

1230 

1900 

5. 


1774 

Mercury 

13.6 

-38.85 

357 

95.78 

.0197 

Prehist. 

Nickel 

8.75 

1452 


7.8 

.142 

1751 

Platinum 

21.37 

1755 

3910 

10. 

.166 

1741 

Potassium 

.87 

62.3 

712 

7.2 

— 

1807 

Silver 

10.5 

960.5 

1955 

1.59 

.990 

Prehist. 

Sodium 

.97 

97.5 

750 


— 

1807 

Tin 

7.29 

232 

2270 

11.5 

.155 

Prehist. 

Tungsten 

18.7 

3400 

5830 

5.6 

.476 

1781 

Zinc 

7.15 

419.4 

930 

5.8 

.265 

1520 

ALLOYS 







Brass 







70Cu:30Zn 

8.6 

900 


7. 

.260 


Bronze 







67Cu:33Sn 

8.6 

750 





Gold coin 







90Au:10Cu 

17.17 

940 





Solder 







67Pb:33Sn 

9.4 

240 


16. 



Steel 







99Fe:lC 

7.7 

1510 


12. 

.115 



Order of Ductility of the Common Metals 


Gold, silver, platinum, iron, copper, zinc, tin, lead. 

Except for iron, the malleability of the metals is also in the above order. 

























80 


GENERAL CHEMISTRY 


(5) Determination of the Melting Point of an Alloy. An 

alloy known as Wood’s metal consists of Bi, 50%; Pb, 25%; 
Sn, 12.5%; Cd, 12.5%. Refer to the table and tabulate the 
melting points of these metals in your notebook. Obtain 40 
to 50 g. of Wood’s metal from the storeroom and determine 
its melting point as follows: 

Place the alloy in an ordinary test tube and clamp this in 
an upright position in a beaker of boiling water. Place a 100° 
thermometer in the Wood’s metal and stir it gently as it melts. 
When the alloy is completely molten and the thermometer reg¬ 
isters about 90°, turn out the gas flame and remove the beaker 
of water. Let the thermometer remain stationary in the melted 
metal and read the temperature every 30 seconds as the alloy 
cools and freezes. Continue the readings for several minutes after 
the alloy has solidified. Record time and temperature in your 
notebook in parallel vertical columns as you take the readings. 
Your temperature readings should descend from about 85° to 
about 60° and will require 10 to 15 minutes. 

Do not attempt to remove the thermometer from the solidi¬ 
fied mass, but melt the latter in the beaker of boiling water as 
before. Remove the thermometer and carefully detach any 
metal which adheres to it. While the alloy is in the molten 
state pour it into a beaker of cold water. If any of the alloy 
remains in the test tube, allow it to cool and scrape it out by 
means of a piece of glass tubing. Return the Wood’s metal to 
a bottle provided for that purpose. 

Construct a form in your notebook as in Figure 23 (graph 
paper may be used) and from your data plot the “ cooling- 
curve” of Wood’s metal. Draw a line through the points 
obtained and indicate the melting (freezing) point of the alloy. 
How does your result compare with the average of the melting 
points of the component metals? 

What is meant by the term, “ Heat of Fusion ”? Explain 
how this is the cause of the “ halt ” in the cooling-curve of a 
molten metal or alloy. 

(c) Amalgams. — Alloys containing mercury* are called 
amalgams. All the common metals except iron form amal¬ 
gams. Contact of the metal with mercury, or contact of a 

* Caution. — All soluble compounds of mercury are extremely poisonous. Mercury 
vapor, if breathed, causes slow poisoning which is cumulative in effect. There is consid¬ 
erable belief that there is danger of slow poisoning from amalgam dental fillings. 


LABORATORY EXERCISES 


81 


more active metal with a solution of a mercury salt will pro¬ 
duce an amalgam. Is there danger of injuring gold or silver 
jewelry when working with metallic mercury? When working 
with mercury salt solutions? If the jewelry contains a con- 



Fig. 23 

siderable proportion of copper would your answer to the second 
question be different? 

Dip a piece of copper or a copper coin into a solution of a 
mercuric salt until it has a gray coating. Remove the copper 
and polish it with a towel or piece of filter paper. Describe 
the result. The mercury will slowly diffuse into the copper. 
Devise a simple method for removing the mercury from the 
copper and, if your instructor approves, proceed. 
























































82 


GENERAL CHEMISTRY 


EXERCISE 65 

THE CHEMICAL PROPERTIES OF THE METALS 

Textbook: 525-528 

(а) General Properties. — Contrast metals and non-metals 
with respect to: 

(1) The action of water on their oxides. 

(2) The action of water on their chlorides. 

(3) The nature of the ions in which they are usually found 
and the pole at which they are liberated in electrolysis. 

Illustrate (1) and (2) by means of equations and (3) by definite 
examples. 

(б) Displacement Series. — By means of appropriate ex¬ 
periments determine the relative order of activity of the 
following metals: 

Zn, Mg, Cu, K, Fe, Hg, Pb. 

For example, if metal A, placed in a solution containing ions 
of metal B, partly dissolves, displacing B, then metal A is 
more active than metal B and belongs above it in the series. 
Assuming both metals to bivalent, the change which occurs 
may be expressed: 

A + B++ = A++ -f B. 

(c) Amphoteric Metals. — Certain elements, usually classi¬ 
fied as metals, are sometimes found in the cation and some¬ 
times in the anion portion of their compounds. Such elements 
are said to be amphoteric, that is, capable of forming both acids 
and bases. Common examples of such metals are: Al, As, Cr, 
Mn, Sb, Sn, Zn. 

Obtain solutions of salts of Al, Sb, Sn, and Zn in which these 
metals form the cations. Use about 3 cc. of each solution in 
separate test tubes and add to each a solution of sodium 
hydroxide, drop by drop, until a precipitate appears. What 
is the precipitate in each case? Write the equations. 

Divide each precipitate into two portions and treat one por¬ 
tion of each with a dilute acid and the other portion of each 
with some more sodium hydroxide solution. In the latter 
cases the hydroxide of the amphoteric metal is acting as an 


LABORATORY EXERCISES 83 

acid toward the stronger base, sodium hydroxide. Write the 
equations. 

Also write formulas for compounds in which As, Cr, and Mn 
form the cations and compounds in which these metals are 
present in the anions. The two types of compounds of the 
seven metals should now be arranged in parallel columns for 
comparison. 

(d) Action with Oxidizing Agents. — The relative ease of 
oxidation of the metals is, in general, the same as the order of 
the metals in the Displacement or Electromotive Series. Any 
apparent exception will probably be due to the nature of the 
oxidation product which, if insoluble, forms a protective coat¬ 
ing on the metal and prevents further action. 

From their relative positions in the Electromotive Series 
would you expect aluminum or zinc to be oxidized more 
readily by nitric acid? Try the experiment, using dilute nitric 
acid, and explain the result. 

Determine whether a ferrous or a ferric salt is formed when 
iron is acted upon by the following: 

1. Dilute hydrochloric acid. 

2. Dilute nitric acid. 

3. Very dilute nitric acid. 

To do this, clean the rust from three small iron nails by immers¬ 
ing them for a moment in a small amount of concentrated 
hydrochloric acid. Wash the nails with water and place them 
separately in three test tubes. Add 5 cc. of dilute hydrochloric 
acid to one tube, 5 cc. of dilute nitric acid to another, and to 
the third test tube add a solution made by diluting 1 cc. of 
dilute nitric acid with 5 cc. of water. Allow the reactions to 
proceed for at least ten minutes. The test tube containing the 
nail and hydrochloric acid should be heated at intervals to 
accelerate the action; do not heat the other two tubes. 

Meanwhile, obtain a solution of a ferrous salt and a solution 
of a ferric salt from the shelf. To each add sodium hydroxide 
and note the character of the precipitate in each case. Equa¬ 
tions. Apply this test to determine whether a ferrous or a 
ferric salt is formed by the reactions occurring in the three test 
tubes. If ordinary dilute nitric acid is made by adding 2 1. of 
water to a liter of the concentrated acid, what is the strength 


84 


GENERAL CHEMISTRY 


of your very dilute nitric acid relative to concentrated nitric 
acid? 

Drop a few iron filings into a test tube which is moist, on 
the inside. The filings should adhere to the walls of the tube. 
Allow this preparation to stand in your desk until the next 
laboratory period, then examine it and answer the following 
questions. When iron oxidizes (rusts) in moist air does the 
product form a protective coating and prevent further action? 
Explain why iron which is to be exposed to the weather is 
usually coated with zinc (galvanized) or tin or is plated with 
nickel. 

By the method already outlined determine whether the oxide 
formed on the iron filings is a ferrous or a ferric compound. The 
rust should be separated from the iron which has not oxidized 
by adding some dilute hydrochloric acid and shaking the tube 
to suspend the oxide particles. This suspension should then be 
decanted from the metal and heated to dissolve the oxide. This 
solution should now be tested to determine whether a ferrous 
or a ferric salt is present. Equations. 

EXERCISE 66 

PROPERTIES OF SODIUM 

Textbook: 538-543; 222-225 

Take a dry watch glass to the storeroom and obtain 3 small 
pieces of sodium metal to use in the following experiments. 

( Caution: The sodium should be handled with forceps and 
must not come in contact with moist fingers nor with water, 
except as directed.) 

(а) General. — Clean the oil from one piece of sodium by 
touching it on all sides with a piece of filter paper. Trim off 
one side of the sodium with a knife or a sharp edge of glass in 
order to expose a fresh surface of the metal. Examine and 
describe the metal with respect to hardness, color, luster, 
tendency to tarnish (oxidize); note its melting point and 
density in the table accompanying Exercise 64. 

(б) Reaction with Water. — Place 10 cc. of water in a small 
beaker, drop into it the piece of sodium examined in (a), and 
quickly cover the beaker with a glass plate. When the reaction 
is complete, test the solution with litmus paper and rub a little 


LABORATORY EXERCISES 


85 


of it between your fingers (?). Equation. Explain how you 
could convert this product into common salt, stating how you 
would determine when you had added the exact amount of the 
required reagent. 

(c) Changes on Exposure to Air. — This experiment will be 
understood only if it is kept in mind that the atmosphere con¬ 
tains carbon dioxide and water vapor as well as oxygen and 
nitrogen. 

Clean the oil from another piece of sodium and scrape off as 
much of the coating as you can. Place the clean metal in a dry 
beaker and cut or mash the sodium into thin pieces in order to 
present a large surface to the air. Leave this exposed in your 
desk, notice it from time to time, and explain the following 
changes which will probably occur during the week: 

1. The preparation becomes a solution. Test it with litmus 
and by touch. What is present? How did it obtain water? 
What is the technical name of this phenomenon? 

2. It becomes crystalline. Test a portion with a dilute acid 
and explain the effervescence. What is the complete formula 
of this crystalline substance? What further change may occur? 

{d) Preparation of Sodium Peroxide. — Place a plug of glass 
wool or cotton in a drying bulb, fill the bulb with granular 
calcium chloride to 
within 5 cm. of the 
end, and plug this end 
also with glass wool or 
cotton. Set up the 
apparatus as shown in 
Figure 24, placing a 
cleaned piece of sodium 
in a small porcelain 
boat in the ignition p IG . 24 

tube. By means of a 

suction attachment draw a slow stream of air through the dry¬ 
ing bulb and over the sodium, which should be carefully heated 
until it melts. If no suction is available, air may be blown 
through from the other end. The molten sodium should take 
fire in the stream of air. Remove the gas flame as soon as the 
sodium begins to glow. When the reaction is complete allow 
the tube to cool, then remove and examine the product. 














86 


GENERAL CHEMISTRY 


To 1 cc. of dilute sulfuric acid add 10 cc. of water, pour this 
into a test tube, and to this add some of the product obtained 
by burning sodium. Test the resulting solution for hydrogen 
peroxide, using 2 cc. of ether, 2 drops of a solution of potassium 
dichromate and shaking (?). What does this prove as to the 
product formed when sodium is burned? Suggest some reasons 
why it is advisable to dry the air before passing it over the 
sodium. What is the action of water on the combustion product 
of sodium? Equation. 


EXERCISE 67 

PREPARATION AND PROPERTIES OF SODIUM 
HYDROXIDE 

Textbook: 225-231 

(a) Preparation. — Write the equation for a reaction by 
which you have previously produced sodium hydroxide. 

Weigh out 15 g. of sodium carbonate on the laboratory scales 
and calculate the weight of quicklime needed to react with this 
weight of material to produce sodium hydroxide. To make this 
calculation it is necessary to know whether the sodium car¬ 
bonate which you are using is anhydrous or a hydrate and, if 
the latter, whether it is the monohydrate or the decahydrate. 
Write the formulas for the three possibilities. Devise a way 
to determine whether or not the sodium carbonate is a hydrate. 
If carefully heated, the decahydrate will melt before decompos¬ 
ing and may thus be distinguished. 

Having calculated the weight of calcium oxide needed, weigh 
out 25 per cent in excess of this amount to allow for impurities. 
Add 15 cc. of water to the lime in a small beaker and allow it 
to stand until the lime has slaked. This will require about 10 
minutes. Equation. In the meantime the sodium carbonate 
should be dissolved in 150 cc. of water in a beaker and heated 
to boiling. Add more water to the slaked lime to make a milky 
suspension and pour about three-fourths of this into the sodium 
carbonate solution. Stir this well and continue to boil it for a 
few minutes. Equation. 

Now allow the precipitate to settle and transfer about 1 cc. 
of the clear solution to a test tube by means of a pipette. Treat 
this with dilute hydrochloric acid and notice whether or not a 


LABORATORY EXERCISES 


87 


gas is evolved (?). If effervescence occurs, all the carbonate 
has not been precipitated and more of the lime must be added. 
Continue the treatment with milk of lime and the testing with 
hydrochloric acid until all the carbonate has been removed from 
solution as CaC0 3 . Allow the precipitate to settle and decant 
the clear solution through a filter paper into a clean flask or 
bottle. Bring all the precipitate on the filter paper, allow it to 
drain and preserve it for use in Exercise 69. 

(6) Tests. — Test small portions of the filtrate with two dif¬ 
ferent indicators to show that it contains the OH - ion. Place 
two or three drops of the filtrate on a watch glass and evaporate 
to dryness. Examine and briefly describe solid sodium hydrox¬ 
ide. Obtain a platinum wire and heat one end of it in the 
Bunsen flame until the wire no longer colors the flame. Now 
touch the red-hot end of the wire to the substance on the watch 
glass and again place the end of the wire in the flame (?). Test 
two other sodium compounds from the side shelf in this man¬ 
ner (?). This characteristic flame color may be used as a test 
for sodium compounds. 

(c) Reaction with Aluminum. — Place a piece of aluminum 
in a test tube and cover it with some of the solution of sodium 
hydroxide. Does any reaction occur? Should aluminum uten¬ 
sils be washed or scoured with materials containing a strong 
alkali? Heat the test tube and contents and determine the 
identity of the gas evolved. To do this it may be necessary 
to collect a test tube full of the gas by the displacement of 
water. Write the equation for the reaction noting that the 
same compound of aluminum is formed here as in Exercise 
65 (c). 

(d) Effect on Textiles. — In separate test tubes place a small 
piece of wool and a bit of cotton cloth. To each add 5 cc. of the 
solution of sodium hydroxide, heat both tubes and note the 
effect. State how you could distinguish between cotton and 
wool by means of a strong alkali. 

Save the remainder of the sodium hydroxide solution for the 
next exercise. 

(e) Analysis of a Wool-Cotton Mixture. — Obtain a piece of 
the mixed cloth about 8 cm. square and carefully weigh it. 
Place this in a beaker and cover it with a 20 per cent solution of 
KOH or NaOH. To prepare this solution dissolve 5 g. of the 
solid alkali in the calculated amount of water. Boil the solution 


88 


GENERAL CHEMISTRY 


for about three minutes, or until the wool is all dissolved. Avoid 
spattering the hot alkali on your flesh or clothing. Remove the 
cotton residue, taking care that no detached threads are lost. 
Rinse thoroughly, make slightly acid with dilute HC1 and rinse 
again. Press the moisture from the cotton and dry it at about 
100° C. by placing it in a dry beaker or dish and heating it 
carefully to avoid scorching. Weigh the residue and calculate 
the percentage composition of the fabric. 


EXERCISE 68 

HYDROXIDES OF THE METALS 

(a) Insoluble and Amphoteric Hydroxides. — Arrange 10 
clean test tubes in the rack and in these place separately about 
2 cc. of solutions containing the following ions: Cu ++ , Mn ++ , 
Fe++, Fe +++ , Zn++, A1+++, Pb++, Ni++, Co++ Cr+++. To each 
of these tubes add about 1 cc. (10 drops) of the sodium hydroxide 
solution prepared in Exercise 60 or the similar solution on the 
shelf. Write the equations, carefully recording the color of 
the precipitate in each case. 

Heat test tube Number 1, containing the copper compound, 
and explain the change which occurs. Carefully note the colors 
first obtained in tubes 2 and 3 (repeat if necessary) and explain 
the changes which occur in these two on standing several 
minutes. 

Now add an excess of sodium hydroxide to each of the ten 
tubes, with shaking, and note whether or not the precipitate 
dissolves. If solution occurs, write the equation — recall 
Exercise 65 (c). 

(i b ) Unstable Hydroxides. — In separate test tubes study as 
in (a) the action of sodium hydroxide on solutions of a silver 
salt, a mercurous salt, and a mercuric salt. Describe the pre¬ 
cipitates. They are the oxides of the metals — recall the change 
in copper hydroxide when heated. Note the position in the 
Electromotive Series of the metals which do not form stable 
hydroxides. 

(c) Classification. — The soluble hydroxides are compounds 
of the metals high in the Electromotive Series. The solubility 
of these hydroxides in water at ordinary temperature is in the 


LABORATORY EXERCISES 


89 


same order as the activity of the metals. That is, KOH is the 
most soluble of the common hydroxides; ammonium hydroxide 
falls in third place. From these facts and your results in parts 
(a) and ( b ) of this exercise, tabulate the hydroxides in your 
notebook as follows: 


Soluble 

Hydroxides 

Hydroxides 
Precipitated by 
NaOH, insoluble 
in Excess 

Hydroxides 
Precipitated by 
NaOH, soluble 
in Excess 
(Amphoteric) 

Oxides 

Precipitated by 
NaOH, Hydroxides 
Unstable 

Most soluble 
1 . 

2. 

3. 

4. 

5. 

6. 

Least soluble 

Formula, Color 

Formula, Color 

Formula, Color 


EXERCISE 69 

SODIUM BICARBONATE AND SODIUM CARBONATE 

Textbook: 543-549 

(a) The Ammonia-Soda or Solvay Process. — Saturate 20 cc. 
of concentrated ammonium hydroxide with sodium chloride in 
a test tube by adding an excess of the salt, shaking the mixture 
vigorously at intervals for some time, and then filtering out the 
undissolved solid. Catch the filtrate in a test tube and bubble 
carbon dioxide through it by means of a delivery tube reaching 
nearly to the bottom of the test tube. The carbon dioxide 
may be prepared by treating the precipitated calcium carbonate 
saved from Exercise 67 (a) with dilute hydrochloric acid in a 
generator as shown in Figure 25. The precipitate from 67 (a), 
if dry, should be broken into small lumps and sufficient water 
should be added to the generator to cover the end of the thistle 
tube. Add dilute hydrochloric acid through the thistle tube, a 
little at a time, as required to keep a slow stream of gas bubbling 
through the solution in the test tube. Continue the experiment 










90 


GENERAL CHEMISTRY 



Fig. 25 


until a considerable amount of precipitate appears in the test 
tube. If this does not occur before the calcium carbonate in the 
generator is exhausted, the latter may be replenished with 
marble chips. If solid carbon dioxide, “dry ice,” is available it 
may be added directly to the solution to produce the bicarbonate. 

Filter the precipitate obtained in the test tube, spread out 
the filter and allow it to dry until it has lost the odor of ammonia. 

The drying may be hastened by placing the 
paper in a dish on a water bath. Examine and 
describe the product, testing it for sodium and 
for the carbonate radical. 

In writing the equations for the formation of 
the precipitate it may be assumed that carbon 
dioxide reacts with ammonium hydroxide to 
form ammonium bicarbonate. This product may 
then react by double decomposition with sodium 
chloride precipitating sodium bicarbonate which 
is the least soluble of the substances involved. 
(6) Preparation of Sodium Carbonate. — In a dry test tube 
place about 2 g. of sodium bicarbonate and fit the test tube with 
a stopper bearing a delivery tube bent to pass downward when 
the test tube is clamped in an inclined position. Arrange the 
lower end of the delivery tube to dip into 5 cc. of limewater in a 
test tube and then heat the sodium bicarbonate. What evidence 
have you as to two of the products of the decomposition of 
sodium bicarbonate? Equation. Is this reaction reversible? 
State exactly how you would proceed in converting sodium 
carbonate into sodium bicarbonate and write the equation. 

(c) Effects with Indicators. — Dissolve about 0.5 g. of sodium 
bicarbonate in 20 cc. of water, divide the solution equally be¬ 
tween two test tubes, add 2 drops of phenolphthalein to one 
and 2 drops of methyl orange to the other. From a pipette 
add dilute hydrochloric, drop by drop, to each tube until the 
indicator changes, noting the number of drops of acid required 
in each case. Does sodium bicarbonate affect one indicator 
more than the other? Does this show that various indicators 
change color at a slightly different H-ion concentration? 

Dissolve about 4 g. of sodium carbonate in a sufficient 
quantity of water to make 100 cc. of solution. Place this solu¬ 
tion in a clean flask or bottle. In another flask place 20 cc. 
of dilute hydrochloric acid and 220 cc. of water and mix well. 











LABORATORY EXERCISES 


91 


If ordinary dilute hydrochloric acid has a strength which is 
6N, what is the normality of the acid solution which you have 
just prepared? 

Obtain a burette, rinse it with some of the acid you have 
prepared and fill it with this acid. Using 25-cc. portions of the 
sodium carbonate solution for each experiment, titrate the 
sodium carbonate solution with the acid solution, using methyl 
orange as the indicator in one sample and phenolphthalein as 
the indicator in a separate sample (?). 

Repeat the titration with each indicator, this time heating 
each 25-cc. sample of sodium carbonate solution to boiling and 
continuing to boil it while making the titration until the color 
change is permanent. Compare the results obtained in the four 
titrations and draw conclusions as to the reaction which occurred 
in each case. Equations. 

EXERCISE 70 

PREPARATION OF POTASSIUM NITRATE 

Textbook: 554-555 

Weigh out 25 g. of sodium nitrate and the amount of potas¬ 
sium chloride required to react with it. Place these two salts 
in a beaker, add 50 cc. of water and heat to boiling. The mix¬ 
ture should be stirred occasionally to hasten the process of 
solution. Filter the hot solution and allow the filtrate to cool. 
If desired, it may be placed in a beaker of cold water which 
should be renewed as it becomes warm. 

From the solubility data in the accompanying table plot the 
solubility curves for the four salts over the temperature range 
0° -100° C. These four curves should appear on a single chart 
(see illustration in Exercise 64) with Temperature on the hori¬ 
zontal axis and Solubility (g. per 100 g. water) on the vertical axis. 


SOLUBILITY IN 100 PARTS OF WATER 


Temperature 

0° 

to 

o 

o 

O 

O 

60° 

80° 

100° 

KC1 

28.5 

34.7 

40.1 

45.5 

51. 

56.6 

kno 3 

13.3 

31.2 

64. 

111. 

172. 

247. 

NaCl 

35.7 

36. 

36.6 

37.3 

38.5 

39.8 

NaNOs 

72.9 

87.5 

102. 

122. 

148. 

180. 












92 


GENERAL CHEMISTRY 


When the solution has cooled to about 20° and crystals have 
formed, decant the mother liquor into another beaker, com¬ 
pletely draining the crystals by inverting the beaker containing 
them over the other beaker as in Figure 26. 

According to the solubility curves you have drawn, which 
one of the four salts must these crystals be? Calculate the 
weight of this salt you would expect to obtain at 20° from the 
amounts of solutes and solvent you have used. Examine and 
describe these crystals as to form (drawing) and taste. In 
reference below these will be called “ A ” crystals. Would these 
crystals be as pure if the solution had been cooled to 0°? Why? 

Evaporate the mother liquor to about one-half its volume, 
stirring it to prevent “ bumping,” and note the formation of 
crystals in the hot solution. These will be 
referred to as “ B ” crystals. Immediately 
decant the hot solution into another beaker 
and allow it to cool. Drain the crystals as 
before. Wash the “B ” crystals by decanta¬ 
tion using about 5 cc. of hot distilled water. 
Examine and describe these crystals as to 
form (drawing) and taste. Make a flame 
test with some of these crystals (?). Dis¬ 
solve a few of these “B” crystals in water 
and test half of the solution for a chloride (?). 
Try the brown-ring test for a nitrate on the 
other half of the solution to determine whether or not the 
crystals have been sufficiently washed. 

When the hot solution has cooled, another crop of “A” 
crystals will be obtained. Drain these as before and combine 
them with the first crop. Dissolve all of the “A” crystals by 
heating them with a small amount of distilled water, 8 to 10 cc., 
and allow them to recrystallize. This process of recrystallization 
tends to purify the crystals. When the solution has cooled 
decant the liquid and bring the crystals into a filter paper 
fitted to a funnel. Allow them to drain and rinse them twice 
with 1-cc. portions of distilled water. 

Dissolve a few of the crystals in water and make the nitrate 
and chloride tests on separate portions (?). If the chloride test 
is positive explain how you could further purify the crystals. 
Allow the crystals to dry on the filter paper and save them for 
Exercise 71. 






LABORATORY EXERCISES 


93 


EXERCISE 71 

USES OF SOME POTASSIUM COMPOUNDS 

Textbook: 556; 549-550 

(a) Preparation of Fuse Paper. — Using the crystals of po¬ 
tassium nitrate prepared in Exercise 70 make a saturated solu¬ 
tion at room temperature by shaking an excess of the salt in a 
test tube with 10 cc. of water. Soak a piece of filter paper and 
a short piece of wrapping cord in this solution and allow them 
to dry in the air. Heat a piece of glass tubing, draw it to a 
capillary and, using this as a pipette, trace a continuous design 
with the solution on a piece of filter paper starting at the edge 
of the paper. Mark the starting point with a pencil and when 
the paper has dried ignite it at this point only. A flameless 
combustion should travel around the design. What purpose 
does the potassium nitrate serve? When the string you have 
prepared is dry, ignite it and compare its rate of burning with 
that of an untreated piece. Twist pieces of the prepared filter 
paper into fuses for use as directed below. 

(&) Potassium Bitartrate — Baking Powder. — Using the re¬ 
mainder of the potassium nitrate solution from (a) warm it to 
about 40° C. and saturate it at this temperature. To 5 cc. of 
this solution add an equal volume of a solution of tartaric acid, 
mix well and allow the tube to stand or hasten the crystalliza¬ 
tion by cooling with water. The precipitate is potassium acid 
tartrate, KHC 4 H 4 0 6 , which is only slightly soluble. It is called 
“ cream of tartar.” Write the equation. Filter out the precipi¬ 
tate, transfer it to a porcelain dish and apply heat, gently at 
first to dry it, then strongly. The tendency to char when heated 
or when treated with concentrated H 2 S0 4 is characteristic of the 
salts of organic acids. When the residue is cool treat it with a 
dilute acid (?). From this test and from the appearance of the 
residue try to formulate the probable equation for the decom¬ 
position of KHC 4 H 4 06 by heat. 

Calculate the weight of sodium bicarbonate required to react 
with 1 g. of cream of tartar to produce KNaC 4 H 4 0 6 . Name this 
salt. Mix the required amounts of the two dry substances with 
a gram of starch. This mixture represents the more expensive 
class of baking powders. Divide the mixture between two dry 
test tubes, to one portion add cold water and to the other add 


94 


GENERAL CHEMISTRY 


boiling water (?). Write the equation. Why did the reaction 
not occur when the salts were dry? Why is starch present in 
commercial baking powders? 

(c) Bengal Fires. — In a clean dry mortar grind 5 g. of po¬ 
tassium chlorate to a fine powder. Pour this on a piece of paper, 
putting aside a small spoonful for (d) below. Wipe out the mor¬ 
tar and pulverize 1 g. of strontium nitrate or chloride. Add this 
to the potassium chlorate and mix these two powders on a piece 
of paper with 1 g. of “ flowers of sulfur.” Mix thoroughly, using 
a spatula or the fingers, but do not grind together. ( Caution: 
Do not grind sulfur or any other combustible material with 
potassium chlorate as a dangerous explosion will result.) Pour 
half of the mixture into a conical pile on a piece of asbestos and 
place a piece of fuse paper, at least an inch long, in the top of 
the pile. Light the end of the fuse paper and stand at a safe 
distance. What is the probable reaction between sulfur and 
potassium chlorate? The salts of strontium give a red color to 
flames. Salts of other metals could be substituted to give 
other colors. 

( d ) Flashlight Powder. — Mix the spoonful of powdered 
KC10 3 saved from (c) above with about an equal amount of 
magnesium powder. Do not grind together. Place this mixture 
in a pile on a piece of asbestos and ignite it using a fairly long 
piece of fuse paper. What is the application of this light? Why? 

EXERCISE 72 
AMMONIUM COMPOUNDS 

Textbook: 558-561 

The ammonium compounds closely resemble the correspond¬ 
ing compounds of potassium in crystal form, general appear¬ 
ance and solubility (as shown in Exercise 73). A few distinctive 
properties are outlined below. 

(a) Ammonium Hydroxide — Repression of Ionization. — 

Dilute 2 cc. of ammonium hydroxide with water to 20 cc., add 
a few drops of phenolphthalein solution and divide the colored 
solution between two test tubes. Save one portion as a standard 
and to the other add 5 g. of solid ammonium chloride and shake 
the tube until the salt is dissolved. Explain the result, writing 
ionic equations to show the equilibria involved and the effect 

pj V W* 0 -> ^ H A// * * & 

Ay. _ ( * r 

it * t t CX <- * .• V. 


^y\ <■ 

r\ 




SUU3L * o 0*10 **-*Uy* *MuKy '*• 

^ /#% ^ . ' 

'V' '—'*> v **4^ {M4 -2} V\ it-a •+ Msf 

LABORATORY EXERCISES ' 95 


of the common ion. Could a similar result be obtained using 
the corresponding sodium or potassium compounds instead of 
NH 4 OH and NH4CI? Why? %d>^ 

To 5 cc. of a solution of a magnesium salt add 2 cc. of am¬ 
monium hydroxide (?). Equation. Repeat the experiment 
using ammonium hydroxide to which an equal volume of am¬ 
monium chloride solution has been added (?). Explain the 
different results. Add some ammonium chloride, either solid 
or in solution, to the test tube containing the precipitate and 
shake it (?). 4MjUU& £CAM£ 

(i b ) Ammonia Complexes. — Precipitate a small amount of 
silver chloride in a test tube and dissolve the precipitate by 
adding ammonium hydroxide. A soluble complex is formed 
having the composition Ag(NH 3 ) 2 Cl. Test the effect of dilute 
nitric acid on this solution (?). Write the equations. Ammonia 
forms similar complexes with some other metals such as copper, 
zinc, cobalt, nickel. The number of molecules of ammonia con¬ 
tained in the complex ion is usually twice the maximum valence 
of the metal. 

To a solution of copper sulfate add ammonium hydroxide 
slowly with shaking until the precipitate which forms is redis¬ 
solved. The color of this complex is sometimes used as a test 
for copper. Write the formula for the copper-ammonia complex, 
j (c) Removal of Ammonium Salts. — In an analysis it is often 
necessary to remove ammonium salts as these interfere with 
certain tests. In a small, dry Pyrex beaker place a “ pinch ” of 
ammonium chloride. Heat the beaker on a wire gauze until all 
the salt has volatilized from the bottom of the beaker (?). The 
decomposition of this salt may be represented: 

NH 4 C1 + heat <=* NH 3 + HC1 






Does this indicate that a sublimate of the salt will condense on 
the cooler portions of the beaker? A reversible decomposition 
is called dissociation. When produced by heat, as in this case, 
it is called thermal dissociation. What name is given to the 
dissociation which occurs when an acid, base, or salt is dis¬ 
solved in water? 

Clean the beaker and repeat the experiment using ammonium 
nitrate instead of ammonium chloride (?). Write the equation 
for the thermal decomposition of this salt. Is the reaction re¬ 
versible? Should we call this dissociation? With what reagent 


96 


GENERAL CHEMISTRY 


might the NH 4 C1 have been treated before heating to effect a 
complete removal of ammonium salt? Show by means of an 
equilibrium equation why your answer is true. 

(, d ) Test for the Ammonium Radical. —- Recall or test the 
action of a strong alkali on ammonium salts and write an 
equilibrium equation to show why this reaction may be used 
to identify ammonium salts. State several means by which 
you can recognize ammonia gas. 

EXERCISE 73 

TESTS FOR POTASSIUM, SODIUM, AMMONIUM SALTS 

Textbook: 561-563 

Most salts of these cations are quite soluble. A few, however, 
are only slightly soluble and may be precipitated if present in 
sufficient concentration — recall KHC 4 H 4 0 6 as a slightly soluble 
potassium salt. 

(a) Perchlorates. — In separate test tubes place 2-cc. por¬ 
tions of solutions of a potassium salt, a sodium salt, and an 
ammonium salt. To each solution add a few drops of perchloric 
acid (storeroom) and describe the results. Equations. How 
could you separate sodium and potassium from a mixture of their 
salts? Alcohol (C 2 H 5 OH) is usually added to decrease the 
solubility of potassium perchlorate and make the separation 
more complete. Why is it necessary to remove ammonium 
salts before testing for potassium with perchloric acid? Recall 
Exercise 72 (c) for the method. 

(i b ) Fluosilicates. — Again using 2-cc. portions of solutions 
of the three salts as in (a) add an equal volume of alcohol (store¬ 
room) and then a few drops of fluosilicic acid, H 2 SiF 6 (store¬ 
room) to each test tube. Allow the tubes to stand for a few 
minutes and carefully note the differences in the nature of the 
precipitates if more than one is obtained. Equations. 

(c) Cobaltinitrites. — To solutions of potassium, sodium and 
ammonium salts add a solution of sodium cobaltinitrite, 
Na 3 Co(N0 2 ) 6 . This reagent is prepared as follows: Dissolve 
1 g. of sodium nitrite in 5 cc. of water, add 2 cc. of cobalt 
nitrate solution and 1 cc. of acetic acid. Shake this mixture, 
allow it to stand a few minutes, and then dilute it with an equal 
volume of water. Use this to make the tests as directed above 


LABORATORY EXERCISES 


97 


and describe the results. Equations. Is it necessary to remove 
ammonium salts before testing for potassium with this reagent? 

(d) Flame Tests. — Clean a platinum wire by repeatedly dip¬ 
ping it in some HC1 solution in a watch glass and heating it in 
the non-luminous flame until it gives practically no color. Test 
several potassium salts, both solids and solutions, with the plati¬ 
num wire in the non-luminous flame (?). Recall or repeat the 
flame test for sodium. Now mix solutions of KC1 and NaCl in 
equal parts and make a flame test with this mixture. Which metal 
gives the predominant color? Is the test for the other metal 
entirely masked? Now repeat the flame test with the mixture, 
this time viewing the flame through a piece of cobalt glass (?). 

Clean the platinum wire and make a flame test with an 
ammonium salt. This dull coloration is given by many sub¬ 
stances and hence there is no characteristic flame for ammonium 
salts. Because of the wide distribution of sodium salts be 
careful to distinguish the brilliant and persistent coloration 
given the flame by considerable amounts of sodium from the 
flash given by a trace such as may be obtained by drawing a 
clean platinum wire between the fingers. 

Tabulate the tests for these three cations somewhat as follows: 



HC10 4 

H 2 SiF 6 

Na 3 Co(N0 2 )6 

Flame 

Na 

K 

nh 4 






( e ) Analysis of Unknown No. 1 . — This unknown will con¬ 
sist of a salt of sodium, potassium or ammonium or a mixture 
of any two or all three. Report only the cations in this analysis. 


EXERCISE 74 

SOAP 

Textbook: 4.99-500 

(a) Preparation of Soap. — “Soap” is the common name for 
the sodium salts of the fatty acids, usually palmitic, stearic or 
oleic. It is made by treating a fat with an alkali, the composi- 








98 


GENERAL CHEMISTRY 


tion of the soap depending upon the nature of the fat which is 
used, for example: 

(Ci 7 H35C02)3C 3 H5 + 3NaOH 3Ci7H 3 5C0 2 Na + C 3 H 5 (OH) 3 
Glyceryl stearate Sodium stearate Glycerine 

(Fat) (Soap) 

This type of reaction is called saponification. Glyceryl stearate 
is the principal fat in beef tallow. In olive oil the fat is glyceryl 
oleate. Lard is a mixture of the glyceryl esters (organic salts) 
of the three acids mentioned above. What then is the composi¬ 
tion of Castile soap made from olive oil? Of soap made from 
hog fat? 

Weigh 15 g. of lard or other fat into a small beaker; allow 
5 g. of sodium hydroxide to dissolve in 5 cc. of water, add this 
to the fat, and place this beaker in another beaker of boiling 
water. Heat the mixture in this way for about an hour, stirring 
occasionally and adding a small quantity of water as needed. 
The contents of the beaker should stiffen. Remove a small 
portion of the mass on a stirring rod and test it with a drop of 
phenolphthalein. Should a good toilet soap give this test? 
Why? What other impurity is present in this soap? Soap is 
fairly soluble in water but is insoluble in a saturated solution of 
sodium chloride. Your product should be purified as follows: 
Remove the mass to a larger beaker, add 100 cc. of distilled 
water and boil this mixture, breaking up any lumps with a 
spatula. Saturate the water with finely powdered sodium chlo¬ 
ride while heating and stirring the solution. Allow the mixture 
to cool and remove the cake of soap. It should be pressed into 
a filter paper fitted to a funnel and allowed to dry there. Test 
its ability to form a lather by shaking a small quantity with 
10 cc. of distilled water. Finally weigh your product and sub¬ 
mit it to your instructor for inspection. 

(6) Detergent Action. — The detergent (cleansing) action of 
soap is dependent upon its property of taking grease and dirt 
into colloidal suspension and thus mechanically removing them 
by emulsification and peptization. 

Prepare a 1 per cent soap solution by shaking 2 g. of soap 
with the calculated volume of distilled water in a flask, heating 
it to complete the solution. Use this solution throughout this 
exercise and save the remainder for Exercises 75 and 77. 

In separate test tubes place 20 cc. of water and 20 cc. of 


LABORATORY EXERCISES 


99 


soap solution. To each tube add 1 g. of animal charcoal and 
shake the tubes vigorously. Compare the amounts of solid 
held in suspension after the tubes have stood for several min¬ 
utes. The action of soap in removing grease and oil is indicated 
in Exercise 75. 

(c) Action with Ca++, Mg++, K+, etc. — Add some soap solu¬ 
tion to a solution of calcium chloride and filter off the pre¬ 
cipitate. What is it? Assuming the soap to be sodium stearate 
write the equation. Shake some of the precipitate with water. 
Does it dissolve? Lather? Does it have the cleansing property 
of a soap? 

Add soap solution to solutions of KC1, MgCl 2 , FeCL and 
any other chlorides. What other metal might be substituted 
for sodium in soap? Soaps of this metal do not harden; they 
are liquids. 

Add 2 cc. of dilute sulfuric acid to 10 cc. of soap solution (?). 
Equation. Shake the tube vigorously to coagulate the pre¬ 
cipitate. What is it? Will it dissolve when boiled with a solu¬ 
tion of NaOH? Equation. 

EXERCISE 75 
EMULSIONS 

Textbook: 148 

A dispersion of droplets of one liquid in another liquid is 
called an emulsion. It is, of course, necessary that the two 
liquids be mutually insoluble (immiscible). It is that type of 
colloidal solution in which both components are liquids. A 
familiar example of an emulsion is that formed from water and 
an oil. 

(a) Preparation of an Emulsion. — Pour 10 cc. of kerosene 
and 10 cc. of water into a test tube. Likewise mix 10 cc. of 
kerosene with 10 cc. of soap solution. Shake both tubes vig¬ 
orously and compare them after they have stood a few mo¬ 
ments (?). Are oil and water miscible? Can a stable emulsion 
be prepared by shaking oil and water together? Explain in 
your own words what happens when an emulsion “ breaks ” or 
separates into layers of its two components. An emulsifying 
agent retards or prevents breaking because it forms a film 
around the dispersed droplets and keeps them from contact and 


100 


GENERAL CHEMISTRY 


union with each other. Does your experiment indicate that 
soap is an emulsifying agent? Save the tubes and contents 
for (6). 

Make a solution of albumin (white of egg) by shaking a 
small quantity of dried albumin with 10 cc. of water in a test 
tube. To this solution add an oil, a few drops at a time, 
shaking the tube vigorously, until 5 cc. of the oil have been 
added (?). 

An emulsion used for salad dressing is made by stirring to¬ 
gether vinegar (dilute acetic acid), olive oil (cottonseed or corn 
oil may be substituted), and eggs. What is the familiar name 
for this emulsion? What is the emulsifying agent in this 
case? 

(i b ) Types of Emulsion. — Make a simple sketch to rep¬ 
resent a highly magnified section of an emulsion. Label the 
droplets “internal phase” or “dispersed phase” and the 
surrounding liquid “external phase,” “continuous phase,” or 
“dispersion medium.” It is obvious that, theoretically, in the 
case of oil and water either may be the internal phase. That 
is, we may have oil droplets dispersed in water or water dis¬ 
persed in oil. It is easy to determine which condition is actual 
by adding more of one of the components to the emulsion. If 
the external phase is added it will mix readily; if the internal 
phase is added it will not mix. 

Shake the tube reserved from (a) containing kerosene and 
soap solution and divide the emulsion into two nearly equal 
portions. To the test tube containing one portion add water 
and to the test tube containing the other about the same 
amount of kerosene. Cover each tube with a thumb and in¬ 
vert it once to mix the materials but do not shake (?). Is 
water or oil the external phase in this emulsion? 

Prepare some magnesium stearate by adding 10 cc. of a 
solution of MgCL or MgSCh to 10 cc. of soap solution. Filter, 
wash, and partially dry the precipitate by pressing it between 
filter paper. Add it to the oil-water mixture saved from (a) 
and shake vigorously (?). 

Divide this material into two equal portions and determine 
as before the identity of the internal and external phases. 
Is magnesium stearate an emulsifying agent? How does the 
action of magnesium stearate differ from that of sodium 
stearate? 


LABORATORY EXERCISES 


101 


EXERCISE 76 

CUPROUS COMPOUNDS 

Textbook: 575-577: 503-506 

(а) Cuprous Chloride. —Dissolve 3g. of cupric chloride in 15cc. 
of water, add 2 g. of copper turnings and 3 cc. of concentrated 
hydrochloric acid. Boil the solution gently for several minutes 
or until the green tint is no longer perceptible. Equation. 

Boil 20 cc. of water in a test tube to remove the dissolved 
oxygen, pour about 1 cc. of the cuprous chloride solution into 
this water (?). What is the precipitate? If a blue color appears 
in the solution was the reduction of cupric chloride complete? 

To 5 cc. of the cuprous chloride solution add a solution of 
sodium hydroxide until the mixture is alkaline (?). Equation. 
Compare this result with that obtained when sodium hydroxide 
solution is added to a solution of a cupric salt. 

Pour the remainder of the cuprous chloride solution into a 
large volume of water (?). Equation. Decant or filter the 
solution and test the solubility of portions of the precipitate in 
dilute hydrochloric acid and in ammonium hydroxide. Expose 
a portion of the precipitate to the air on a piece of filter paper (?). 
Explain your results. 

(б) Cuprous Iodide. — Dilute 3 cc. of copper sulfate solution 
with an equal volume of water and add 3 cc. of potassium 
iodide solution (?). Filter out the precipitate and account for 
the color of the filtrate (?). Dilute some of the filtrate and 
test it with starch solution (?). Write the equation for the 
action of a cupric salt with potassium iodide. 

(c) Fehling’s Solution. — Dissolve 3.5 g. of copper sulfate 
crystals in water and dilute the solution to 50 cc. Place this 
in a small bottle or flask and label it “ Solution A.” Dissolve 
5 g. of solid sodium hydroxide and 12.5 g. of Rochelle salt (po¬ 
tassium sodium tartrate, KNaC4H 4 06-4H 2 0) in water together 
and dilute this solution to 50 cc. Bottle this and label it 
“ Solution B.” 

Mix 5 cc. of Solution A with 5 cc. of Solution B (?). What 
would have been the result in the absence of the Rochelle salt? 
To this mixture add 1 cc. of glucose solution and boil the solu¬ 
tion noting all changes. What is the precipitate? What was 


102 


GENERAL CHEMISTRY 


the action of glucose on the cupric salt? This reaction is used 
in detecting and estimating glucose and other sugars of the 
same type. 

Repeat the Fehling test using a solution of cane sugar in¬ 
stead of glucose (?). “ Invert ” some cane sugar by boiling 

10 cc. of its solution with a few drops of concentrated hydro¬ 
chloric acid for several minutes. Carefully neutralize the acid 
with sodium hydroxide solution and test this invert sugar with 
Fehling’s solution as above — that is, mix 5 cc. of Solution A 
and 5 cc. of Solution B; add 2 or 3 cc. of the solution to be 
tested and boil several minutes (?). Hydrogen ions catalyze 
the hydrolysis of cane sugar (C 12 H 22 O 11 ) to form glucose and 
fructose, both of which have the formula C 6 H 12 O 6 . Write the 
equation for this “ inversion.” 

Prepare a starch solution by moistening a pinch of starch 
with water and adding boiling water to this. Test some of this 
starch solution with Fehling’s solution as directed above (?). 
To 2 cc. of thin starch solution in a test tube add about 5 cc. 
of saliva, warm the tube slightly by immersing it in hot water 
and shake it for several minutes. Test the resulting solution 
with Fehling’s solution (?). The first stage in the digestion of 
starch is its conversion to sugars through the agency of the 
enzyme diastase in the saliva. 

(d) Testing Common Substances for Glucose. — If any of 

the following substances are available test them for glucose (or 
similar sugars): Table syrups such as corn syrup, molasses, 
maple syrup; fruit juices — orange, apple, raisin, etc.; jams 
and jellies; candies. Use a clear solution of the material for 
the test. 


EXERCISE 77 

PREPARATION OF SOME CUPRIC COMPOUNDS 

Textbook: 575-578 

(a) CuS0 4 • 5H 2 0. — Weigh 5 g. of copper, place this in a 
clean beaker, and add 10 cc. of concentrated sulfuric acid, 20 cc. 
of dilute nitric acid and 20 cc. of water. Heat this mixture over 
a low flame in the hood until the copper has all dissolved, adding 
more dilute nitric acid if necessary. Add more water if crystals 


LABORATORY EXERCISES 


103 


begin to form before all the copper has dissolved. While this is 
in progress proceed with (6) of this exercise. When solution is 
complete, filter the liquid if it is not clear, catch the filtrate in a 
clean beaker and evaporate this solution until crystals begin to 
form. Place the beaker in cold water and stir the liquid until 
crystallization is complete. Decant the liquid from the crystals, 
evaporate the mother liquor and crystallize it as before. Com¬ 
bine the two crops of crystals, drain them and dissolve them in 
the minimum amount of boiling water in a small beaker. Cover 
this solution with a watch glass and set it in your desk for a day 
or more to recrystallize. 

Decant the liquid from this last crop of crystals, rinse them 
with a small amount of distilled water, and dry them between 
filter papers. Weigh the crystals and calculate your per cent 
yield from the weight of copper used in the experiment. Select 
a good crystal, rinse it again, dissolve it in water and test small 
portions for: (1) Sulfate; (2) Nitrate (as impurity); (3) Copper 
(use the ammonium hydroxide test and also test with a solution 
of potassium ferrocyanide). Copper gives a chocolate-colored 
precipitate of cupric ferrocyanide with this reagent; (4) Also 
test a small dry crystal of your product for water of hydration. 
What is the color of the anhydrous salt? 

Assuming that nitric acid acted as the oxidizing agent, being 
reduced to nitric oxide, write the equation for the production 
of copper sulfate in this experiment. 

Ask your instructor to inspect the copper sulfate crystals and 
your record for the experiment. Save the crystals for (c). 

( b ) Cu(NH 3 ) 4 S0 4 .H 2 0. — Grind 10 g. of copper sulfate to a 
fine powder, place this in a small flask and add 20 cc. of con¬ 
centrated ammonia solution. Shake the flask for some time and 
then warm gently to complete the solution of the solid. Trans¬ 
fer the solution to a beaker, cover it and place it in your desk 
to crystallize. Filter off the crystals and dry them very 
carefully with filter paper. Submit the crystals to your in¬ 
structor for inspection. 

(c) CuS0 4 - (NH 4 ) 2 S0 4 -6H 2 0. — Using the crystals of blue 
vitrol prepared in (a) or some from the reagent shelf, grind 
10 g. to a fine powder. Calculate the weight of ammonium 
sulfate equivalent to the 10 g. of copper sulfate. Weigh out 
this amount of ammonium sulfate, grind it with the copper 
sulfate in the mortar, transfer the mixture to a small flask or 


104 


GENERAL CHEMISTRY 


beaker and dissolve it in the smallest possible amount of boiling 
water. Acidify this solution with a few drops of dilute sulfuric 
acid and immediately filter it while hot into a small beaker. 
Cover the beaker and set it in your desk for a day or more to 
crystallize. 

Drain, rinse and dry the crystals, and submit them to your 
instructor for inspection. Compare the crystals of the three 
substances you have prepared in this exercise and state how 
you could distinguish them from each other, either by appear¬ 
ance or chemical tests. Clearly distinguish by name between 
the compound prepared in (6) and that prepared in ( c) of this 
exercise (?). 

EXERCISE 78 

METALLIC COUPLES 

Textbook: 578 

(a) Displacement. — Place a piece of zinc and a piece of 
copper (not in contact) in the bottom of a beaker and cover 
them with dilute sulfuric acid. What reaction occurs? 

Bring the two metals into contact (?). Is the rate of the 
reaction changed? Explain the evolution of gas from the sur¬ 
face of the copper. How could you determine whether or not 
any copper is dissolved? Remove the piece of copper and try 
the effect of touching the zinc under the acid with a platinum 
wire (?). Is any energy evolved in this reaction? In what form 
does it appear? Explain how the addition of copper sulfate 
accelerates the production of hydrogen from zinc and sulfuric 
acid. 

Pour a few drops of lead nitrate solution into a watch glass 
and place a piece of zinc in this for a few moments (?). Equa¬ 
tion. The zinc should be only partially submerged in this solu¬ 
tion. Amalgamate a piece of zinc by covering it with a solution 
of a mercuric salt for a few moments. Equation. Place these 
two pieces of zinc in separate test tubes and in a third tube place 
an ordinary piece of zinc. Cover each preparation with dilute 
sulfuric acid, observe and record what happens. In a zinc 
amalgam the mercury dissolves in the zinc exposing a surface 
which is equivalent to pure zinc. Ordinary zinc contains other 
metals as impurities. Does pure or impure zinc react more 
readily with an acid? Explain. 


LABORATORY EXERCISES 


105 


From the standpoint of the E.M.F. series, which of two 
metals in contact will dissolve first in an acid? From the sur¬ 
face of which metal is most of the hydrogen liberated? To 
answer this question cover a piece of amalgamated zinc (pure 
zinc) with dilute sulfuric acid and touch the zinc under the 
liquid with a platinum wire (?). 

(b) Action of Metals with Water. — Boil some water in a 
beaker to remove the dissolved air. In separate test tubes place 
a piece of magnesium ribbon, a piece of aluminum and a piece 
of zinc. Add some of the boiling-hot water to each of these 
samples and observe whether or not any hydrogen is produced. 
Now add a few drops of a solution of a mercuric salt to the 
test tubes containing magnesium and aluminum and some 
copper sulfate solution to the test tube containing zinc. Again 
heat the water to the boiling point and notice if any of the 
metals displace hydrogen. Explain. 

EXERCISE 79 

ELECTROCHEMICAL ACTION 

Textbook: 578-581 

(Two students may work together on this exercise) 

(a) Electricity from Chemical Action. — Obtain two zinc 
electrodes and two copper electrodes from the storeroom. 
Place one zinc electrode and one copper electrode in a beaker 
and hold these in place and apart by slipping a smaller beaker 
within the other. Mix some dilute sulfuric acid with twice its 
volume of water and pour some of this solution into the outside 
beaker. Add water to the inner beaker if necessary to give it 
weight. Is any action visible at either electrode? Connect the 
free ends of the wires leading to the electrodes. State what 
happens at each electrode while the cell is thus short-circuited. 
Disconnect the wires to break the electrical circuit while the 
cell is not in use. 

Amalgamate the other zinc electrode, rinse it, and set up a 
separate cell similar to the one just described. Observe this 
cell for action at the electrodes when the circuit is open and 
closed, and compare this cell with the one first prepared. What 
is the purpose of amalgamating the zinc electrode? 


106 


GENERAL CHEMISTRY 


Connect the two cells in series by joining the copper electrode 
of one to the zinc electrode of the other. Determine the sign of 
each electrode (+ or -) by means of the following polarity 
tests: Add about 3 cc. of thin starch solution to 5 cc. of potas¬ 
sium iodide solution. Moisten a piece of filter paper with some 
of this mixture and save the remainder for (6). Touch the ends 
of the free wires to the moistened paper — the ends of the wires 
should be about 1 cm. apart. This test may also be performed 
by pouring some of the solution in a watch glass and holding 
the wires 1 cm. apart in the liquid (?). What causes the color? 
At which electrode does it appear? A similar polarity test may 
be made with a solution of sodium chloride containing a 
few drops of phenolphthalein. Test as before and explain the 
appearance of the color. At which electrode is the color seen 
in this case? 

Determine the voltage (E.M.F.) of this battery of two cells 
by connecting its wires to the proper terminals of a voltmeter (?). 
Write an ionic equation to show the source of the electric current 
in these cells. 

(6) The Daniell Cell. — Rinse the zinc and copper electrodes 
used in (a), amalgamate the other zinc electrode, and obtain 

two porous cups from the store¬ 
room. Assemble two cells as 
illustrated in Figure 27, placing 
dilute sulfuric acid in the po¬ 
rous cup and a saturated solu¬ 
tion of copper sulfate in the 
beaker surrounding the porous 
cup. Is any chemical action 
visible? Is any action visible 
when the circuit is completed 
by joining the wires from the electrodes? Make a polarity test 
as directed in (a). This test will be obtained more readily if the 
cells are first connected in series. Using the voltmeter as before, 
determine the E.M.F. of a single Daniell cell, the E.M.F. of the 
two cells connected in series, and the E.M.F. of the two cells 
connected in parallel. Compare these values (?). When the 
two cells are connected in series sufficient current may be ob¬ 
tained from them to cause a small flashlight bulb to glow. 

(c) Electrolysis. — Dip the free ends of the wires from the 
Daniell cells (connected in series) into a solution of copper sul- 












LABORATORY EXERCISES 


107 


fate. Does electrolysis occur? What happens at each pole? 
What practical application is illustrated by this experiment? 

Rinse the ends of the wires and dip them into a solution 
of sodium hydroxide. What occurs at each electrode in this 
solution? Explain how this experiment could be used as a 
polarity test. What would have been the products of this 
electrolysis if the electrodes had been platinum instead of 
copper? 

Remove some of the liquid from one of the porous cups by 
means of a pipette and test it for zinc ions as follows: Make 
the solution alkaline with ammonium hydroxide and pass in 
hydrogen sulfide. Refer to Exercise 84 (d), first paragraph. A 
white precipitate, soluble in hydrochloric acid, is zinc sulfide. 
What occurs at each electrode of the Daniell cell when it is in 
operation? Express the net chemical change in the cell by 
means of an ionic equation and indicate the source of the 
electrical energy. 


EXERCISE 80 

SILVER 

Textbook: 581-590 


AjjrU^ 



(a) Displacement of Silver. — Place a drop of metallic mer¬ 
cury in 5 cc. of silver nitrate solution in a test tube and allow 
this preparation to stand undisturbed until the next laboratory 
period. At that time describe the result and write the equation. 

Mix 2 cc. of silver nitrate solution and 2 cc. of cupric sulfate 
solution in a test tube. Dilute this mixture with an equal 
volume of water and drop into it a piece of zinc (?). Describe 
the deposit which appears immediately and that which appears 
after a few minutes. Is the black deposit copper? To decide 
this drop a piece of zinc into a mixture of 2 cc. of silver nitrate 
and 6 cc. of water and compare the results (?). After the solu¬ 
tion which contained both copper and silver has been in con¬ 
tact with the piece of zinc for at least 10 minutes, remove about 
2cc. by means of a pipette. Test j^arate P or li° ns this f° r 
' * ^copper, see Exercise 77 (a), and for silver. Devise a test for 
silver based on the fact that silver chloride is insoluble in water 
and dilute acids but soluble in ammonium hydroxide. Is copper 




108 


GENERAL CHEMISTRY 


or silver completely displaced from the solution? Leave the re¬ 
mainder of the solution in contact with the zinc until the next 
laboratory period and test the solution for copper, silver and 
zinc (?). Write the equations. 

(b) Effect of Light. — Using silver nitrate solution as an ink, 
mark some design on a piece of filter paper. Allow this to dry 
and then warm it gently (?). The same effect is produced by 
longer exposure to light. Mark a towel or other cloth with 
silver nitrate solution and notice it after a day or two (?). The 
stain will not wash out. Can you suggest a use for this ink? 
Similar stains will be made on the skin by silver nitrate solution. 
Why is silver nitrate usually kept in a brown bottle? Note 
the stains at the mouth of the bottle of silver nitrate solution. 
What is this stain? How would you prove it? 

^Wd3T5~cc7of sodium chloride solution to 5 cc. of silver nitrate 
solution, filter quickly and transfer half of the precipitate to 
another filter paper. Immediately place one portion of the 
precipitate in the dark and expose the other portion to the light. 
Notice the changes which occur in the latter and finally com¬ 
pare it with the other portion which was left in the dark. Scrape 
the surface of the portion exposed to light (?). Has the action 
penetrated to a considerable depth? The effect of light is to 
make the silver salts more readily reducible to metallic silver. 
Explain the practical application of this fact. What is a “de¬ 
veloper”? Mention one. What is a “fixing agent ”? Mention 
the one commonly used. 

(c) Reactions of the Silver Ion. — Add sodium hydroxide 
Solution to 2 cc. of silver nitrate solution (?). Refer to Exer¬ 
cise 68 (6). Equation. What other cations give similar results? 

Add ammonium hydroxide, drop by drop, to 1 cc. of silver 
nitrate solution (?). Write two equations to illustrate your 
observations. What molecules and ions are present in the solu¬ 
tion called ammonium hydroxide? Show this by indicating the 
equilibria involved. Which ion or molecule is responsible for 
the first change noted in the reaction with silver nitrate? Which 
causes the second change? What complex ion is finally present? 
Refer to Exercise 72 (6). 

Pour 2 cc. of silver nitrate solution into each of three test 
tubes and dilute each portion with about 5 cc. of distilled water. 
By adding solutions of the proper reagents precipitate silver 
chloride in one of these tubes, silver bromide in the second, and 

iO x t -7 dUetfM'y 





G-<s.X '"ft 

i /***.„„* 

LABORATORY EXERCISES 109 

silver iodide in the third. Divide each precipitate into two 
portions and test the solubility of the three silver halides in 
ammonium hydroxide and in a solution of sodium thiosulfate. 
Tabulate your results so that it is obvious how these three 
silver salts may be distinguished by color and by solubility. 

( d ) Silvering a Mirror. — Carefully clean a watch glass by 
washing it with a hot, 10 per cent solution of sodium hydroxide, 
rinse it first with tap water and finally with distilled water. 

Pour 4 cc. of silver nitrate solution into the watch glass and add 
ammonium hydroxide, drop by drop, until the precipitate 
which first forms has dissolved. Add about 3 cc. of a 10 per 
cent glucose solution to the solution in the watch glass and 
warm the mixture by placing the watch glass over a beaker 
of boiling water. Continue heating the water bath until a con¬ 
siderable deposit of silver is produced. What chemical prop¬ 
erty of glucose is utilized here? Recall its action in Exercise 
76(c). 

Rinse the watch glass and dry it by pressing a clean cloth or 
filter paper against it. Examine the mirror from the under 
side of the watch glass. Write an equation to show how this 
deposit may be removed from the glass. 


EXERCISE 81 


♦ r <v 

Exerc 


REACTIONS OF CALCIUM, STRONTIUM AND BARIUM 

Textbook: 597-617 > M ^ 1 

(M Flame Tests. — Clean a platinum wire as directed in 
Exercise 73 (d) and test a solution of calcium chloride in the 
flame (?). Repeat the test witW strontium and barium salts, 
being careful to clean the wire after each test. The flame given 
by lithium salts is not easily distinguished from the strontium 
flame; copper compounds and boric acid give a flame color 
similar to that of barium salts. It is not easy to distinguish 
calcium, strontium and barium by means of the flame test 
when two or three of these metals are present. For this 
reason additional tests must be employed in detecting these 
cations. 

(&f Precipitation Tests. — Perform the following tests sep¬ 
arately on 5-cc. portions of solutions of calcium, strontium and 


lo C-C 


A_ 


A A 



barium salts. Describe the results, write the equations , and 
collect your data in a table similar to the form suggested below. 
Indicate the form and color of the precipitates. 


1. Add a solution qf sodium carbonate or ammonium car- 




j *<*) 


n 



bonate (?). 

2. Add a solution of ammonium oxalate (?). 

3. Add a solution of potassium dichromate (?). 1 

4. Add a saturated solution of calcium sulfate (?). 


- a 




4 


! .y' y] 

In test 4 note whether the precipitate forms quickly or only 
after warming and standing. From this observation draw a 
conclusion as to the relative solubility of the three sulfates. 
Of the carbonates and oxalates, the calcium salts are least 
soluble. Indicate the order of solubility for the sulfates, car¬ 
bonates, oxalates and chromates of the alkaline earths, remem¬ 
bering that the strontium compound is always intermediate in 
solubility. 



Carbonate 

Oxalate 

Chromate 

Sulfate 

Flame 

Calcium 

Strontium 

Barium 

wt li/Suf* ■ 

V 

> ft 

k 






(c) Separation. — Starting with a solution containing Ca ++ , 
Sr++, and Ba ++ (use chlorides or nitrates) separate the barium 
as a precipitate (?). Filter out the precipitate and divide the 
filtrate into two portions. In one portion precipitate strontium 
by adding a saturated solution of calcium sulfate and warm¬ 
ing (?). To the other portion add a solution of ammonium 
sulfate, boil and filter. Test for calcium in this filtrate by add¬ 
ing a solution of ammonium oxalate (?). 

(d) Analysis of Unknown No. 2. — This unknown will con¬ 
sist of a salt of calcium, strontium or barium or a mixture of 
any two or all three. Report only the cations in this analysis. 










LABORATORY EXERCISES 


111 


EXERCISE 82 


HARD WATER 


Textbook: 607-609 


t O «->€- 

(a) Soft Water. — To 2fitec. of distilled water in a test tube 
add soap solution, a few drops at a time from the 10-cc. grad¬ 
uate cylinder, shaking the tube vigorously, to determine the ~ 
volume of soap ^o^tfeon required to give a collar of lather or 
suds which will last at least 10 seconds. Record this value for 
comparison below. Repeat this test with tap water and de- 
termine if it has any degree of hardness by comparing the 
amount of soap needed to give a lather in this case with the 
value obtained for distilled water. \ 


(6) Temporary Hardness. — Generate carbon dioxide in a 
” flask, using marble chips and dilute sulfuric acid, and 



jCSs this gas into 20 cc. of limewater until the precipitate which 
forms at first has completely dissolved. What was the pre¬ 
cipitate? What is now present in the solution? Equations. 
Using 10 cc. of this solution determine the volume of soap 


solution required to give a lather (?). Boil the other portion of 
the solution for some time and note the change which occurs (?) . 


UUIUUIUII ^ v/v ^--—■■ - 

the solution for some time and note the change wmcn occurs 


Equation. Now make the soap test on this boiled solution and 
compare the volume of soap solution required with that used 
in the portiotf no^oiled (?). What compounds cause tempo¬ 
rary hardness in water? How may temporarily hard water be 
softened? Explain how it would be possible to soften such water 
by adding exactly the required amount of limewater? Equation. 

(c) Permanent Hardness. — Determine the volume of soap 
solution required to give a lather with 2D cc. of a saturated 
solution of calcium sulfate solution and compare this value 
with that obtained for soft water (?). To 20*cc. of the calcium 
sulfate solution add a gram of solid sodium carbonate and mix 
well by shaking the tube (?). Equation. Mgke the soap test 
on this solution and compare with the values above (?). What 
has been accomplished by the use of sodium carbonate? Re¬ 
peat the test using borax in place of sodium Carbonate (?). 
What compounds may be present in water of permanent 
hardness? What is the practical distinction between the two 
types of hard water? Explain the white deposit formed in 
water kettles and boilers in limestone regions, i 







112 


GENERAL CHEMISTRY 


EXERCISE 83 
MAGNESIUM 

Textbook: 622-627 

(a) Properties of the Metal. — Hold a piece of magnesium 
ribbon by one end with the forceps and heat it about 2 cm. 
from the other end by means of a lighted match. Is there any 
evidence of melting? Now heat the ribbon at the end to ignite 
it, using the Bunsen flame if necessary, and hold the burning 
ribbon over a small beaker to catch the solid product. Assum¬ 
ing that the burning magnesium combines to some extent with 
both of the principal components of the atmosphere, write two 
equations to show the products formed. Add about 2 cc. of 
water to the products in the beaker, heat the water to boiling 
and test the solution with litmus paper (?) . Write the equations 
for the action of water upon the two combustion products. What 
use is made of the fact that the light from burning magnesium 
is rich in ultra-violet rays? Refer to Exercise 71 (d). 

Write equations for the action of magnesium with the dilute 
solutions of the three common acids on the desk. 

(b) Some Magnesium Salts. — Place a few crystals of mag¬ 
nesium chloride on a watch glass and expose it to the atmos¬ 
phere for a day or more (?). Compare this result with the action 
of a calcium chloride under the same conditions. 

Test concentrated solutions of magnesium chloride and mag¬ 
nesium sulfate with litmus and explain the result. To make this 
test, crystals of the salt may be placed on moistened litmus 
paper. 

Place a gram of magnesium chloride crystals, MgCl 2 • 6H 2 0, 
in a porcelain crucible and heat it. Test the escaping va¬ 
por with moist litmus paper (?). Explain. Finally heat the 
crucible at a red heat until the vapor is no longer evolved. 
What is the residue? Equation. Mention a practical use of 
this substance. 

Heat a small quantity of magnesite in a hard glass test tube. 
and pass the gas which is driven off into some lime water (?). 
Equations. Heat the magnesite until the decomposition is 
complete, empty the residue into a small beaker and add 


LABORATORY EXERCISES 


113 


about 10 cc. of water. Boil the water and test the solution with 
fitmus paper and a portion of it with phenolphthalein (?). 
What is present in the solution? Did all the residue dissolve? 
Suggest a reagent which will dissolve it readily, and test your 
idea. Equation. 

(c) Precipitation of Magnesium — Test. — Refer to Exer¬ 
cise 72 (a), second paragraph, for the effect of ammonium salts 
in preventing the precipitation of magnesium hydroxide. Re¬ 
peat the experiment if necessary. 

To 5 oc. of a solution of magnesium salt add a solution of 
ammonium carbonate (?). Equation. To 5 cc. of the mag¬ 
nesium salt solution add an equal volume of ammonium 
chloride solution, mix by shaking the test tube, and then add 
ammonium carbonate solution (?). Explain. 

Repeat the last paragraph using a solution of a calcium, 
strontium or barium salt in place of the magnesium salt (?). 
On this basis explain how you could separate the alkaline 
earth elements from magnesium. 

To 5 cc. of a solution of a magnesium salt add an equal volume 
of ammonium chloride solution, then make the solution strongly 
alkaline with ammonium hydroxide and add a solution of sodium 
phosphate (?). Why was the ammonium chloride added? The 
precipitate is MgNH 4 P0 4 . Describe its appearance, name it, 
and write the equation. This is used as a test for magnesium. 
State how you would use the test conversely for a phosphate. 
What should be the composition of “ magnesia mixture ” used 
in testing for phosphate? Filter out the precipitate, transfer 
some of it to a porcelain crucible and heat this strongly. Test 
the escaping vapor with moist litmus paper (?). Write the 
equation for the thermal decomposition of the salt and name 
the residue. In quantitative analysis phosphorus and mag¬ 
nesium are usually determined by weighing them in this form. 

(d) Separation of Ca++ and Mg ++ . — Dissolve 0.5 g. of dolo¬ 
mite (formula ?) in dilute hydrochloric acid, heating to com¬ 
plete the reaction. Equation. A few particles will probably 
fail to dissolve. These will be silica, an impurity in the dolo¬ 
mite. Dilute the solution to about 20 cc. and filter out the 
insoluble particles. Dissolve a gram of solid ammonium chlo¬ 
ride in the filtrate, make it distinctly alkaline with ammonium 
hydroxide, and add about 10 cc. of ammonium carbonate solu¬ 
tion (?). Equation. Filter out the precipitate and add a solu- 


V 




o 


114 


GENERAL CHEMISTRY 


tion of sodium phosphate to the filtrate (?). For what ion is this 
a test? Equation. Dissolve some of the precipitate on the filter 
paper by pouring acetic acid on it, catch this filtrate in a clean 
test tube and add ammonium oxalate solution (?). Equation. 


EXERCISE 84 




ZINC AND CADMIUM 

Textbook: 


jJtik 


631-636, 


V JU 




^ (a) Basicity. — Test solutions of a zinc salt and a cadmium 
salt with litmus paper ( 7)7 Explain. Name two other metals 
whose salts give a similar reaction. 

^ (b) Hydroxides. — Dilute 5 cc. of a solution containing zinc 
ions with an equal volume of water and add a solution of 
hydroxide, drop by drop, until a considerable pre¬ 
cipitate is obtained. Describe it and write the equation. Di¬ 
vide the precipi£gte hjto two portions, treat one with an excess 
of sodium hydroxide, and the other with dilute hydrochloric 
acid (?). Indicate the manner in which zinc hydroxide may 
ionize. [Refer to Exercise 68 and mention several other hydrox¬ 
ides which may act in this manner.] 

Repeat the directions in the preceding paragraph, substitut- 


. 


ing cadmium for zinc. Compare the results with those obtained 

n . Ai\A/JXJb 


■ u ' 




for zinc. 


Hf 




{m j 


(c) Ammonia-Complex Ions. — Add ammonium hydroxide, 




zmp' 


by drop at first and then in excess, to a solution q 
salt (?). Equations. Repeat this test using a cadmium sa 
solution and compare the results. Write the formulas for 
similar complexes formed by two other metals. 

(d) Sulfides. — In this and subsequent experiments requir¬ 
ing the use of hydrogen sulfide gas, use an Erlenmeyer flask 
fitted with a one-hole stopper bearing a short piece of glass 
tubing to which about 15 cm. of small rubber tubing is attached. 
See Figure 28. The solution to be treated with H 2 S is placed 
in this flask, the rubber tubing is connected to the hydrogen 
sulfide jet and the gas is turned on. The stopper should be 
loosened momentarily to allow some of the air to escape from 
the flask as it is displaced by H 2 S. Hydrogen sulfide is quite 
poisonous and caution should be observed in its use. The 
method just described is designed to avoid the escape of the 





LABORATORY EXERCISES 115 

gas into the atmosphere of the laboratory. Prepare a flask as 
described above and keep it for the purpose indicated. 

Dilute 10 cc. of a zinc salt solution with 20 cc. of water and 
saturate this with hydrogen sulfide gas as directed above (?). 
Now add about 15 cc. of ammonium hydroxide and again treat 
with hydrogen sulfide (?). Explain. Shake the flask and pour 
10-cc. portions of the suspension into each of two test tubes. 
In one test the solubility of the precipitate in acetic acid (?). 
To the other portion add dilute hydrochloric acid, a few drops 
at a time, shaking the tube after each addition, 
until the solution is just neutral and then until it 
just clears (?). Record the number of drops re¬ 
quired to dissolve the precipitate after the solu¬ 
tion has been neutralized. Now add some sodium 
acetate solution until the precipitate reappears. 

Explain how sodium acetate can increase the con¬ 
centration of the sulfide ion in this solution and 
indicate the ionic equilibria which are involved. 

Dilute 10 cc. of cadmium salt solution with 
20 cc. of water and saturate the solution with hy¬ 
drogen sulfide. Compare the appearance and amount of this 
precipitate with the zinc sulfide obtained under the same condi¬ 
tions. Shake up this precipitate and test the solubility of portions 
of it in acetic and dilute hydrochloric acids. The hydrochloric 
acid should be added a few drops at a time as suggested in the 
previous paragraph. Record the number of drops of acid re¬ 
quired. Compare the solubilities of the two sulfides prepared. 

( e ) Summary. — Tabulate the products formed when zinc 
and cadmium ions react with sodium hydroxide, ammonium 
hydroxide, and hydrogen sulfide. Indicate likewise the physical 
and chemical properties of these products. How could you 
separate zinc hydroxide from cadmium hydroxide? How could 
you separate the sulfides of these two metals? 

EXERCISE 85 
MERCURY 

Textbook: 637-6^2 

^(a) Mercurous Salts. — Add son^hydrochloric acid to a 
solution of mercurous nitrate (Tf?xVhffti other cation that you 
have studied would give a similar result? Allow the precipitate 



Fig. 28 




MUil 


AJkt 


* J 

116 GENERAL CHEMISTRY ;'VV, /<T 

to settle, pour off most of the liquid, and add ammonium hydrox¬ 
ide (?). Refer to your textbook for the composition of the prod¬ 
uct. Equation. This is a test for mercurous salts. 

To small portions gf^ mercurous nitrate solution add sep¬ 
arately (1) sodiunf hydroxide /solution; (2) potassiym iodide 
solution; (3) ammonium'laydro^e. Describe the results and 
write the equations. * ^®' ** 

Dilute 2 or 3 cc. of a mercurous salt solution and saturate it 
with hydrogen sulfide gas if). In this and in certain other re¬ 
actions mercurous salts react as if they consisted of a mercuric 
salt and mercury. The products in this reaction are, hence, 
mercuric sulfide and mercury. Equation. 

L s(b) Mercuric Salts. — Add a few drops of hydrochloric acid 
to a solution of a mercuric salt. Compare the solubilities of 
the two chlorides of mercury. What are the common names of 
these two substances? Mention a use for each. 

Add a solution of sodium hydroxide to a mercuric salt solu¬ 
tion. What is the com^Sfeifion of the product? What other 
cations give a similar result? Is the precipitate soluble in 

excess sodium hydroxide? Compare the corresponding results 
„ , . , . (uubvv i 

for cadmium and zmc. 

Add a solution of potassium iodide to a solution of a mercuric 
salt, first adding a small quantity and then an excess (?). 


_ _ x # What 

is finally formed? 

To a solution of mercuric chloride add ammonium hydroxide 
until a definite result is obtained (?). Compare this with the 
corresponding reactions of zinc and cadmium. A solution of 
mercuric nitrate^may be used for this test if the excess nitric 
acid in the solution is first carefully neutralized by adding the 
proper amount of sodium hydroxide solution before adding/ 
ammonium hydroxide. Perform this experiment and write 
the equations. Now add a solution of ammonium nitrate to 
the white precipitate (?). Is the reaction reversible? 

Add a solution of stannous chloride to a solution of mercuric 
chloride, first a small amount and then an excess (?). Notice 
two distinct changes and write the equations. This test is used 
to identify mercuric salts. 

rWXJ Dilute 5 cc. of a mercuric salt solution and precipitate mer¬ 
curic sulfide from this solution. Filter out the precipitate, 
place portions of it in porcelain dishes and test its solubility 
separately in hydrochloric and nitric acids and in aqua regia (?). 


- 4 * 








nri 




Yl H. 


'-/l t 

V 


t 7 



LABORATORY EXERCISES 


117 


Heat these reagents to boiling when making the tests. When 
solution is accomplished evaporate it almost to dryness, take 
up the residue with water and test this solution by adding 
stannous chloride (?). Equations. 

(c) Preparation of Mercuric and Mercurous Nitrates. — 
Place a small drop of mercury in a clean beaker and add 3 cc. 
of dilute nitric acid. Add one or two drops of concentrated 
nitric acid to start the chemical action. Allow this to proceed for 
several minutes, and then decant the solution into a test tube, 
retaining the remainder of the mercury in the beaker. Test 
the solution for mercurous and mercuric ions (?). Equations. 

Add sufficient concentrated nitric acid to the remaining mer¬ 
cury to dissolve it completely and heat the beaker to accelerate 
the action. When the mercury has entirely disappeared test 
the solution for mercurous and mercuric ions (?). Equations. 
What would occur if metallic mercury were placed in the bottle 
containing mercuric nitrate solution? 

EXERCISE 86 

THE PREPARATION OF MERCURIC THIOCYANATE 

Textbook: 61$ 

( Pharaoh’s Serpent Eggs ) 

Carefully counterpoise a small beaker on the balances, and 
in it weigh (to an accuracy of 0.1 gram) about 15 g. of either 
mercuric oxide or metallic mercury. To the oxide in the 
beaker add 15 cc. of concentrated nitric acid, or 20 cc. of con¬ 
centrated nitric acid if metallic mercury is used, and warm 
gently until solution is complete. Evaporate the solution 
slowly over a low flame until a few crystals begin to appear on 
the surface of the liquid. What are they? Write the equation. 

Wash this material into a 400-cc. beaker, using about 200 cc. 
of water altogether, and stir well. If a white suspension or 
cloudiness results, it indicates the formation of a basic salt 
and may be cleared up by adding a few drops of concentrated 
nitric acid and stirring. A faint cloudiness may be disregarded. 
Suggest a formula for this basic salt. 

If a soluble thiocyanate is now added in an amount exactly 
equivalent to the Hg" 1 ^* ions present, insoluble Hg(SCN )2 will 
be precipitated quantitatively. However, an excess of either 


Ill 


GENERAL CHEMISTRY 


of the reagents increases the solubility of Hg(SCN) 2 due to 
the formation of complex ions, and the yield of the precipitated 
product will be lowered proportionately. What is the usual 
effect of an excess of a common ion? In order to avoid an excess 
of either reagent it is necessary to use an indicator which shows 
when the end of the reaction is reached. For this purpose a 
ferric salt may be used since Fe(SCN) 3 is a soluble substance 
and has an intense red color. If a solution of a thiocyanate is 
added to the solution of a mercuric salt containing a few drops 
of a ferric salt solution, the SCN - ions combine with Hg ++ ions 
as long as any of these are present, forming Hg(SCN) 2 which 
appears as a precipitate toward the end of the reaction. When 
the Hg ++ ions are used up and the SCN - ions are added in 
excess, the latter combine with Fe +++ ions to form the red com¬ 
pound, thus indicating that the end of the principal reaction has 
been reached. 

To the mercuric nitrate solution in the beaker add about 
5 drops of ferric chloride solution and then add slowly apd 
with constant stirring a solution of ammonium thiocyanate (or 
potassium thiocyanate) until a faint pink color persists in the 
solution. Equations. 

Allow the precipitate to settle, decant the clear liquid and filter 
the remainder, using a large filter paper (about 18 cm.). Wash 
the precipitate in the filter paper until the washings no longer 
have a pink tint. Allow the precipitate to drain for some time, 
then open the filter paper and lay it flat in a warm place to dry 
or leave it in your desk until the next laboratory day. 

When quite dry, the material should be carefully scraped from 
the filter paper into the original counterpoised beaker and weighed. 
Calculate the theoretical yield and your percentage yield. 

Take several small portions of the material and test its solu¬ 
bility in a solution of mercuric nitrate and in a solution of the 
thiocyanate. 

Mercuric thiocyanate, when made into pellets and ignited, 
burns to a coherent ash which assumes sinuous forms — hence 
called Pharaoh’s serpents. The dried material should be 
moistened to a thick dough with dextrin solution and molded 
with the fingers into conical “eggs.” When these are thor¬ 
oughly dried they may be ignited under the hood on a piece 
of asbestos. ( Caution: Do not breathe the fumes as they 
are poisonous.) 


LABORATORY EXERCISES 


119 


EXERCISE 87 

ALUMINUM 

Textbook: 652-657 

(а) Aluminum Hydroxide. — Dilute 5 cc. of a solution of an 
aluminum salt with an equal volume of water and to this add 
sodium hydroxide solution slowly until a considerable amount 
of precipitate is obtained (?). Equation. Suspend the pre¬ 
cipitate by shaking and divide it into three portions. To one 
portion add an excess of sodium hydr^idfe^sfrlution (?), to the 
second add dilute hyfecfilofic acid (?), and to the third add 
ammonium hydroxide"(?). In which of these reagents is the 
precipitate readily soluble? Equations. Indicate the manner 
in which aluminum hydroxide may ionize. What term is 
applied to hydroxides which have this property? Compare 
zinc, cadmium and magnesium hydroxides with aluminum 
hydroxide, stating how you could distinguish these four sub¬ 
stances from each other by means of reagents mentioned in 
this paragraph. 

Test solutions of aluminum salts with litmqs ^and conclude 
as to the strength of aluminum hydroxide as a base (?). What 
would occur if the aluminum salt of a weak acid were treated 
with water? Would the fact that aluminum hydroxide is 
practically insoluble have any bearing on the result? With 
these answers in mind, add som^ jsodium carbonate solution to 
a solution of an alummumsafeEquation. Add a solution 
of borax to a solution of an aluminum salt (?). Explain. 
SaLujcateT.Q-. c .c. of ammonium hydroxide with hydrogen sulfide ' 
'gas. Equation. f. Add Uu| solution to the solution of an alu¬ 
minum salt (?). Equation! 

Precipitate a small quantity of aluminum- hydroxide by 
means of ammonium hydroxide, filter out the precipitate and 
transfer some of it to a porcelain crucible. Heat this gently at 
first until it is dry and then to a red heat. Allow the crucible 
to cool and examine the residue. What is it? Equation. Is 
the reaction reversible? Test the solubility of the product in 
hydrochloric acid (?). What conclusion could you draw as to 
its melting point? For what is this substance used? 

(б) Baking Powder. — Write the equation for the reaction 
of alum with sodium bicarbonate to liberate carbon dioxide. 


120 


GENERAL CHEMISTRY 


It is suggested that you balance the equation by assuming the 
formation of aluminum bicarbonate whi 



down or is hydrolyzed, 


Calculate the weight of alum needed to react with 2 g. of 
sodium bicarbonate, weigh out these two substances, powder 
the alum, and mix them. Place the mixture in a small 
beaker and add 50 cc. of water (?). After effervescence has 
practically ceased identify the precipitate suspended in the 
solution (?). 

(c) Alum. — Calculate the weight of potassium sulfate 
which will unite with 12 g. of aluminum sulfate crystals, 
A1 2 (S0 4 )3-18H 2 0, to form an alum. Weigh out these quanti¬ 
ties of the two salts, mix them in a beaker and add about 50 cc. 
of water. Heat the water to boiling to dissolve the salts and 
set it aside to crystallize. If a loop of thread is suspended in the 
solution by hanging the ends of the thread over opposite sides 
of the beaker, the crystals will form on the thread in a manner 
resembling rock candy. 

After several days remove some of the crystals, dry them, 
and test them for aluminum, potassium, sulfate, and water of 
hydration. Write the equation for the formation of this alum. 
Mention two metals which would be substituted for aluminum 
in this formula. 

What could be substituted for potassium? Write formulas 
for two other alums. 

Leave some clear alum crystals exposed on a watch glass in 
your desk. 

After several days examine them and discuss any change 
which you observe. 

(d) Alum as a Mordant. — Dissolve about a gram of alum 
or aluminum sulfate in 15 cc. of water. Dilute 15 cc. of methyl 
orange solution with an equal volume of water. Obtain two 
pieces of white cotton cloth about 5 cm. X 10 cm. (Pieces of 
your towel will serve.) Dip one piece of cloth in half of the 
methyl orange solution and allow it to dry. Dip the second 
piece of cloth in the alum solution, squeeze out most of the 
solution, then dip it in the other half of the methyl orange 
solution and allow the cloth to dry. When the pieces of cloth 
have become entirely dry wash them thoroughly with water 
and compare the amounts of dye remaining in the two pieces. 
Explain. 


LABORATORY EXERCISES 


121 


EXERCISE 88 

QUALITATIVE ANALYSIS OF BAKING POWDER 

Unknown No. 3. — This material will be a sample of a com¬ 
mercial baking powder. One type of baking powder was studied 
in Exercise 71 ( b ). Other acidic substances commonly used in 
place of cream of tartar are calcium ac id p hosp hate, potassiu m 
bisulphate, or sodium aluminum sulfate. Write the equations 
by which these substances would react with baking soda — 
refer to Exercise 87 ( b ). Starch is usually present to dilute the 
mixture and to keep it dry. 

Mix about 2 g. of the baking powder with 30 cc. of water and 
identify the gas evolved (?j^%This indicates the presence of 
what anion? When effervescence has ceased, filter the suspension 
and use the residue for (a), reserving the filtrate for other tests. 

(a) Starch. — Puncture the filter and wash the residue into 
a test tube with about 10 cc. of water. Boil this, allow it to 
cool, and test it for starch by adding a drop of iodine solution 
or a solution of potassium iodide to which chlorine water has 
been added. 

( b ) Tartrate. — To 5 cc. of the filtrate add 5 drops of con¬ 
centrated sulfuric acid and evaporate to dryness. A charring | ^ , 
of the residue and the characteristic odor of burnt sugar indi¬ 
cates a tartrate. 

If the baking powder contains “soluble starch’’ this will 
interfere with the above test for a tartrate as it will give the 
same charred appearance. An additional test for a tartrate 
should be made as follows: Shake about a gram of the baking 
powder with 10 cc. of water until action ceases. Filter the solu¬ 
tion, add 2 cc. of sodium hydroxide solution and mix well. 

Now add a solution of copper sulfate, drop by drop, shaking 
the tube, until about 8 drops have been added. If a light blue - ^ 
precipitate of copper hydroxide does not appear the presence 
of a tartrate is indicated. What use was made of this reaction 
in Exercise 76 (c)? QjkjtiXfiU pH'? U> 

(c) Sulfate. — To 5 cc. of the filtrate add barium chloride 
solution. A precipitate insoluble in dilute hydrochloric acid 
indicates a sulfate. Carbonates and phosphates give precipi¬ 
tates which are soluble in acids. 




122 


GENERAL CHEMISTRY 


(i d ) Phosphate. — To 5 cc. of ammonium molybdate solu 
tion add about 1 cc. of the filtrate and acidify the mixture with 
nitric acid. Warm the mixture gently and allow it to stan 
The formation of a yellow precipitate indicates a phosphate 
(e) Aluminum. — Refer to Exercise 87 (6) and decide 

whether most of the aluminum would be retained by the filter 
or would pass through in solution (?). Test some of the filtrate 
for aluminum (?). Devise and perform a test for aluminum 
using some of the original sample. 

(J) Calcium. — Shake a small quantity of the baking powder 
with dilute hydrochloric acid. Filter and neutralize the filtrate 
with ammonium hydroxide. Now make the solution slightly 
acid with acetic acid, heat to boiling and add a solution of 
ammonium oxalate. What result do you expect if calcium is 
present? \M tjU ^ 

(g) Sodium and Potassium. — Moisten a clean platinum 
wire with hydrochloric acid, dip thg wire in the sample of 
baking powder and test it in the flam A. (?). view the flame 
through cobalt glass to test for potassium. If these metals 
are indicated by the flame tests they should be confirmed by 
tests studied in Exercise 73. If aluminum and calcium are 
present these should be removed by precipitation with am¬ 
monium carbonate. The filtrate from this precipitate should 
be evaporated to dryness and the ammonium salts removed (?); 
the residue from this treatment, dissolved in water, may be 
tested for sodium and potassium.^ 

From your results indicate the substances present in the 
sample of baking powder. Report this conclusion as the 
analysis of your unknown. 


EXERCISE 89 


TIN 

Textbook: 666-669 


(a) Stannic Compounds. — Dilute 5 cc. of stannic chloride 
with an equal volume of water and treat the solution with 
hydrogen sulfide (?). Equation. Filter out the precipitate, 
transfer most of it to a porcelain dish and add 3 or 4 cc. of 
yellow ammonium sulfide, (NH^S*. Warm the dish gently 
(do not boil) and stir the mixture until the precipitate dissolves. 


LABORATORY EXERCISES 


123 


What is formed? Equation. Filter this solution into a test 
tube and acidify it with dilute hydrochloric acid (?). Compare 
this result with that obtained when some yellow ammonium 
sulfide is acidified. Equations. This experiment is used as a 
test for stannic compounds. 

( b ) Stannous Compounds. — Repeat the above experiment 
throughout using stannous chloride in place of stannic chloride. 
Compare the result at each point with the corresponding result 
for the stannic compound, and write all the equations. These 
reactions are used to identify stannous compounds. 

Add a few drops of mercuric chloride solution to 2 cc. of 
stannous chloride solution. Note two distinct changes, adding 
more stannous chloride if necessary. Equations. State how 
this experiment may be used as a test for one type of mercury 
compound and conversely for one type of tin compound. 

Add a few drops of potassium dichromate solution to 5 cc. of 
stannous chloride solution (?). Write the equation taking ac¬ 
count of the presence of excess HC1 in the stannous chloride 
solution. Repeat this test using potassium permanganate 
solution instead of the dichromate (?). Equation. What 
seems to be the chief chemical property of stannous chloride? 

To 5 cc. of stannous chloride solution add a solution of 
sodium hydroxide, slowly at first and then in excess (?). Write 
two equations to represent the changes which you observe. 
Name the final product. Add portions of this solution to: (1) a 
solution of bismuth chloride; (2) the white precipitate, BiOCl, 
made by diluting bismuth chloride solution with water; (3) bis¬ 
muth hydroxide, made by adding ammonium hydroxide to a 
solution of bismuth chloride. The black product is metallic 
bismuth. Write the equations. 

(c) Action of Tin with Acids. — Treat a small piece of tin 
..with 5 cc. of concentrated hydrochloric acid in a small beaker. 
Warm the beaker gently to accelerate the action, allow this to 
proceed for some time, and finally evaporate the liquid almost 
to dryness. Allow the beaker to cool. Is there any evidence 
of a crystalline product? Add 20 cc. of water to the beaker to 
dissolve the product, discard any tin which remains undis¬ 
solved and test the solution to determine whether stannous or 
stannic chloride was produced (?). 

Add a few drops of concentrated nitric acid to a small piece 
of tin (?). Describe the product. What is it? How could you 


124 


GENERAL CHEMISTRY 


distinguish tin from zinc and from aluminum by means of 
nitric acid? If you are in doubt try the experiment. 

To a very small piece of tin add 3 cc. of concentrated hydro¬ 
chloric acid and 1 cc. of concentrated nitric acid. What is 
this mixture of acids called? Allow the preparation to stand 
until all the tin has been dissolved, dilute the solution with 
20 cc. of water and determine whether a stannous or a stannic 
compound has been formed (?). Write the equation for the 
solution of tin in this mixture of acids. 

(d) Preparation of Stannic Chloride. — Place 2 g. of tin in a 
dry test tube, immerse this for most of its length in a bottle of 
cold water, and pass chlorine into the test tube until the tin 
has been converted to stannic chloride. The chlorine must be 
dried by passing it through a drying bulb filled with granular 
calcium chloride. Design your apparatus and a suitable 
method for generating the chlorine and obtain the approval of 
your instructor before proceeding with the experiment. 

Describe the stannic chloride you obtain. Determine its 
boiling point as follows: Support the test tube containing the 
stannic chloride with a clamp attached to an iron stand. Wipe 
the tube dry on the outside. Obtain a cork stopper which is 
considerably larger than the mouth of the test tube, bore a hole 
in the cork and insert a thermometer. By resting the cork on 
the top of the test tube, suspend the thermometer in the tube 
so that the bulb of the thermometer is 2 cm. above the surface 
of the liquid. Boil the liquid, using a very low flame, so that the 
vapors do not rise much above the middle of the test tube but 
condense on the sides and run down. When a constant tem¬ 
perature is obtained, record it as the boiling point of the liquid. 

When the stannic chloride is cool, pour a portion of it (with 
care) into some water (?). Suggest a method for recovering 
tin from “ tin plate,” that is, iron plated with tin. 

EXERCISE 90 f M 

LEAD 

Textbook: 673-680 

(a) Test for Lead. — Dilute 5 cc. of lead nitrate with an 
equal quantity of water and add hydrochloric acid until pre¬ 
cipitation is complete (?). Filter out the precipitate and dis- 



LABORATORY EXERCISES 


125 


solve it by pouring successive small portions of boiling water 
on the precipitate in the filter, catching the solution in a beaker. 
To small portions of this hot solution add separately: (1) dilute 
sulfuric acid (?); (2) a solution of potassium dichromate (?); 
(3) hydrogen sulfide solution or gas (?). Write the equations 
and describe the products by which you identify lead. 

( b ) Lead Hydroxide. — Add a solution of sodium hydroxide, 
drop by drop, to 5 cc. of lead nitrate solution until a considerable 
amount of precipitate has been formed (?). Suspend the pre¬ 
cipitate by shaking, divide it into three portions and test its 
solubility in: (1) excess sodium hydroxide solution; (2) am¬ 
monium hydroxide; (3) acetic acid. Write equations to illus¬ 
trate the results obtained. How does lead hydroxide differ in 
behavior from zinc hydroxide? 

(c) Oxides of Lead. — Obtain small samples of lead monox¬ 
ide, lead dioxide, and red lead. Give the formulas, common 
names, and a brief description for these three compounds. Test 
the solubility of each oxide in dilute nitric acid, warming the 
acid if necessary. Does this experiment give any indication of 
the probable constitution of red lead? Explain. 

Treat small quantities of each of the three oxides with con¬ 
centrated hydrochloric acid (?). Equations. Which of these 
oxides oxidize HC1? 

Add bromine water to 5 cc. of sodium hydroxide solution 
until a yellow color is perceptible in the solution. Add this 
solution to 5 cc. of a solution of lead acetate (?). Transfer this 
mixture to a small beaker and boil it for several minutes (?). 
Identify the product by means of appropriate tests and write 
the equations resulting in its formation. 

(d) Separation of Pb ++ , Hg + , and Ag+. — Mix 2 cc. of lead ni¬ 
trate solution, 2 cc. of mercurous nitrate solution, and 2 cc. of sil¬ 
ver nitrate solution in a test tube or small beaker and add dilute 
hydrochloric acid until no further precipitation occurs. Filter out 
the precipitate wash it with 10 cc. of cold water, and treat it on 
the filter paper with small portions of boiling water. Catch the 
solution which runs through and test portions of this for lead by at 
least two methods. Continue the treatment with hot water until 
the washings which run through do not give a test for lead. 

To the residue on the filter add ammonium hydroxide and 
catch the solution which runs through. Acidify the filtrate 
with dilute nitric acid and identify the precipitate (?). Ex- 


126 


GENERAL CHEMISTRY 


amine the residue on the filter. What is it? Recall Exercise 
85 (a). Write the equations for the reactions involved in the 
separation and identification of these three ions. 

(e) Unknown No. 4. — Test this unknown solution for lead, 
silver, and mercurous ions by the procedure outlined in ( d ). 


EXERCISE 91 


CHROMIUM 

Textbook: 688-692 


(a) Chromic Hydroxide. — Precipitate some chromic hydrox¬ 
ide by the addition of either sodium hydroxide solution or am¬ 
monium hydroxide to a solution of a chromic salt. Test the 
solubility of this precipitate in: (1) excess ammonium hydrox¬ 
ide; (2) a dilute acid; (3) excess sodium hydroxide. Write 
equations. Boil the clear solution in (3) until a change 
occurs (?). Transfer about 2 cc. of this suspension to a small 
beaker andadd about a gram of sodium peroxide, a very little at a 
time, with stirring. Add about 10 cc. of water and boil to decom¬ 
pose the excess sodium peroxide. A yellow color denotes the pres¬ 
ence of sodium chromate. Write the equation for its formation. 
To a portion of the solution add a solution of barium chlo¬ 
ride (?). Equation. Save a portion of the solution for (6). 

(b) Chromates and Dichromates. — Acidify the solution re¬ 
served from (a) with dilute nitric acid and note the change in 
color which occurs (?). Equation. Add sodium hydroxide 
solution to a solution of potassium dichromate until a change 
in color is noted (?). Equation. 

Add 2 drops of potassium dichromate solution to 5 cc. of 
water, acidify the solution with 2 drops of dilute sulfuric acid, 
add about 2 cc. of ether and 1 cc. of hydrogen peroxide solution. 
Shake the tube and note the color in the ether layer (?). This 
may be used as a test for chromates and dichromates, or con¬ 
versely for hydrogen peroxide. It is not a test for chromic salts. 

To small portions of potassium dichromate solution add 
separately: (1) a solution of barium chloride; (2) a solution 
of a lead salt; (3) a solution of silver nitrate. The precipitates 
are chromates and may be used as identifying tests — recall 
Exercises 81 (6) and 90 (a). Write the equations for the forma¬ 
tion of the precipitates. Divide the precipitate in (1) into two 


f 


LABORATORY EXERCISES 


127 


portions and test the solubility of barium chromate in acetic 
acid and in dilute hydrochloric acid (?). 

What is the action of a reducing agent on a dichromate? 
Illustrate your answer by means of an equation. 

Outline a test for a chromic salt. 

(c) Chromic Anhydride. — Dissolve 5 g. of sodium dichro¬ 
mate in 10 cc. of water, warming it to hasten the solution. 
This salt is much more soluble than potassium dichromate. 
Cool the solution and add with care about 15 cc. of concentrated 
sulfuric acid. This mixture is used in laboratories as a cleaning 
solution for glassware. Allow the solution to cool slowly and 
to stand for some time. Filter out the crystals using a small 
plug of glass wool in the stem of the funnel. Why may a filter 
paper not be used? 

Examine and describe the crystals. Dissolve a few of them 
in water (?). Add lead nitrate solution to the solution of the 
crystals (?). Pour some sodium hydroxide solution over the 
remainder of the crystals in the funnel and catch the solution 
which filters through. Judging from its color what compound 
is present? Acidify this solution with nitric acid (?). 

(i d ) Preparation of Ammonium Dichromate. — Dissolve 10 g. 
of chromic anhydride in 20 cc. of water and add the calculated 
amount of concentrated ammonium hydroxide (sp. gr. = 0.9; 
28 per cent NH S ). Do not allow the solution to become too 
hot when mixing because much of the ammonia will escape if 
this occurs. Pour the solution into a beaker and allow it to 
crystallize. 

When crystals have formed, filter them out and dry them 
with filter paper. Place the dry crystals in a conical pile on an 
asbestos mat and ignite them by means of a piece of fuse 
paper (?). Write the equations for the formation and decom¬ 
position of ammonium dichromate. 

EXERCISE 92 

PREPARATION OF CHROME ALUM 

Textbook: 690 

Construct the equation for the reduction of potassium di¬ 
chromate by hydrogen sulfide in the presence of sulfuric acid. 
Weigh out 15 g. of potassium dichromate, place this in an 


128 


GENERAL CHEMISTRY 


Erlenmeyer flask and add 100 cc. of water. Calculate the 
volume of concentrated sulfuric acid (sp. gr. = 1.84; 95 per 
cent H2SO4) required for this weight of the dichromate accord¬ 
ing to the equation you have written. Add this volume of 
acid slowly to the water in the flask and heat the mixture to 
boiling to dissolve the potassium dichromate. 

Pass hydrogen sulfide gas into this hot solution in the manner 
described in Exercise 84 (d), shaking the flask vigorously at in¬ 
tervals. This treatment should be continued until the solution in 
the flask is a dark green and will require 30 minutes or longer. 

When the reduction of the dichromate is complete, pour the 
solution into a beaker and boil it for about 5 minutes to coagu¬ 
late the sulfur and expel the excess hydrogen sulfide. Filter 
this solution while hot and set the filtrate aside to cool. De¬ 
scribe the residue on the filter. 

Crystallization of the alum will be facilitated by seeding the 
solution with crystals which may serve as nuclei for growth. 
Since all the alums have the same crystalline form (are iso- 
morphous) crystals of any alum may be used as the seed. When 
the filtered solution of chrome alum is quite cool, drop in a 
piece of ordinary alum about the size of a grain of corn and place 
the solution in your desk to crystallize. 

After several days pour off the mother liquor into another 
beaker and examine the crystals. Compare the color of the 
crystals by transmitted light with the color of the solution (?). 
Pick out the crystal which grew on the seed, break it through 
the center and describe its appearance. Dry a crystal with 
filter paper and leave it exposed to the atmosphere for some 
time. Dissolve some of the chrome alum crystals in water 
(color?) and test portions of the solution for the chromic ion, 
potassium ion and sulfate ion. 

Mixed Crystals. — Prepare 10 cc. of a saturated solution of 
common alum at about 40°. Using 10 cc. of the mother liquor 
from the chrome alum and a few of the chrome alum crystals, 
prepare a similarly saturated solution. Mix these two solutions, 
cool the mixture to room temperature and, if no crystals are 
visible in the solution, seed it with a very small crystal of any 
alum. After sufficient growth has taken place examine and 
describe the crystal comparing it with chrome alum and with 
ordinary alum. What is the chemical composition of this crys¬ 
tal? How could you prove your answer? 


LABORATORY EXERCISES 


129 


EXERCISE 93 

MANGANESE 

(а) Manganous Hydroxide. — Dilute 5 cc. of a manganous 
salt solution with 15 cc. of water, boil this solution to expel 
dissolved oxygen, and divide the solution into four portions. 
To a portion of the solution add ammonium hydroxide (?). 
Is the precipitate soluble in an excess of this reagent? 

To a portion of the solution add an equal volume of ammo¬ 
nium chloride solution and then add ammonium hydroxide (?). 
Recall Exercise 72 (a) and explain the result of this test. 

To the third portion of the manganous salt solution add a 
solution of sodium hydroxide (?). Test the solubility of this 
precipitate in dilute nitric acid (?). 

Add sodium hydroxide to the fourth portion of the solution, 
then add 5 cc. of water and shake the tube to saturate the solu¬ 
tion with air. Note and explain the gradual change in the color 
of the precipitate. Test the solubility of this precipitate in 
dilute nitric acid (?). Now add some hydrogen peroxide solu¬ 
tion and shake the mixture (?). 

Precipitate another small portion of manganous hydroxide 
and treat the precipitate in the alkaline solution with a solution 
of hydrogen peroxide (?). Explain the effect of hydrogen per¬ 
oxide in the two instances, noting its action in acid solution and 
in alkaline solution. Write the equations. 

(б) Manganates and Permanganates. — Make a mixture of 
about 0.5 g. of manganese dioxide with 2 g. of potassium car¬ 
bonate and 2 g. of potassium nitrate. Place this mixture in an 
iron crucible and heat it until all of the material fuses together. 
A green color in the melt indicates the presence of a manganate. 

Allow the crucible to cool and extract the mass with cold 
water, noting the color of the solution. Boil a portion of this 
solution and, if this does not change its color, pass carbon 
dioxide gas into it (?). Write the equation for the formation 
of potassium manganate assuming that the other products 
are carbon dioxide and potassium nitrite. Write an equation 
to explain the change in color of a manganate solution in 
presence of carbon dioxide. 

Add 5 cc. of concentrated nitric acid to a few milligrams of 


130 


GENERAL CHEMISTRY 


freshly prepared manganese dioxide and note that the oxide 
is not soluble. Now add a small amount of lead dioxide, boil 
the mixture for several minutes, add 10 cc. of water and filter 
the mixture. Note and explain the color of the filtrate. The 
products of the reaction are the acid which gives the colored 
ion you observed, lead nitrate, and water. Write the equation. 

EXERCISE 94 

IRON —FERRIC AND FERROUS COMPOUNDS 

Textbook: 722-727 

(a) Tests for Ferric Iron. — Dilute 2 cc. of a solution of ferric 
chloride with 10 cc. of water, divide the resulting solution into 4 
portions and add the following reagents to separate portions: 
(1) sodium hydroxide; (2) potAsfc^ierrocyanide; (3) potas¬ 
sium fdrncyanide; (4) a soluble thiocyanate. Write equations. 

(b) Tests for Ferrous Iron. — Repeat the tests in (a) using a 
diluted solution of a ferrous salt. The thiocyanate test gives 
no color with ferrous compounds. If a red color is obtained in 
this test what is your conclusion as to the action of air upon a 
solution of a ferrous salt? Pour 2 or 3 cc. of ferrous sulfate 
solution into a test tube, add 1 cc. of dilute sulfuric acid and a 
few iron filings. Boil this mixture gently for some time, dilute 
the product with 10 cc. of water, allow the filings to settle to 
the bottom and try the thiocyanate test on this solution (?). 
Draw a conclusion from this test, whatever the result (?). 

Tabulate the results of (a) and (b) indicating the distinguish¬ 
ing colors and the important tests for the two kinds of iron ions. 

(c) Oxidation — Reduction. — Boil 5 cc. of ferrous sulfate solu¬ 
tion with iron filings and dilute sulfuric acid. Decant the solution 
into another test tube, add concentrated nitric acid, drop by drop, 
and boil the solution until further addition of nitric acid does not 
develop a brown color. Note the cofor^oF&e solution (?). Equa¬ 
tion. Test the solution for ferrous and ferrufiron (f). f ° 

Prepare a small amount of ferrous hydroxide, filter it and 
leave the precipitate exposed on the filter paper for 10 minutes. 
Note any change in color (?). Dissolve this precipitate by 
pouring dilute hydrochloric acid on the filter and test the solu¬ 
tion for iron of each valence (?). 

Dilute 2 cc. of a ferric salt solution with 10 cc. of water, heat 
the solution and treat it with hydrogen sulfide gas. Note a 




LABORATORY EXERCISES 


131 


change in color and the formation of a precipitate (?). 
tion. Determine by test the resulting valence of the iror 



(d) Ferrous Sulfide. — precipitate some ferrous sulfide from 

a solution of ferrous Equation/ Describe the pre¬ 

cipitate and test its solubility in dilu?eacid. What use is 
made of ferrous sulfide in chemical laboratories? Given a so¬ 
lution containing AgN0 3 , and Fe(N0 3 ) 2 , outline a 

procedure for the separation and identification of each cation. 

(e) Hydrolysis of Ferric Chloridp. — Test a solution of 
ferric chloride with litmus paperl^ClJonstruct an ionic equi¬ 
librium to explain the result. 

Pour a few drops of a ferric chloride solution^t^ a small 
quantity of boiling water and note the change in color. Refer 
to Exercise 23 and explain. 


EXERCISE 95 
CORROSION OF IRON 

Textbook: 720-722 


Attach a small piece of zinc to the head of a clean iron nail. 
Be sure to obtain good contact between the two metals. Simi¬ 
larly attach a piece of tin to another nail. Place these in sepa¬ 
rate test tubes and in a third tube place an ordinary iron nail. 

Prepare some “ferroxyl reagent ” as follows: Boil about 0.5 g. 
of agar in 50 cc. of water until the agar has been disintegrated 
and taken into colloidal solution. Now add 2 drops of phe- 
nolphthalein solution and a few drops of potassium ferricyanide— 
enough to give a slight yellow tint to the solution. Pour enough 
of this solution into the test tubes to cover the nails, stopper 
the tubes, and set them where they will not be disturbed. 

Examine the tubes at intervals for several days. The purpose 
of the agar was to form a jelly which would hold the materials 
in place. Where the iron goes into solution (rusts) a blue pre¬ 
cipitate of ferrous ferricyanide, Fe 3 [Fe(CN) 6 ] 2 , will be formed. 
Compare the nails as to the amount and location of this action. 
Points in the solution at which hydrogen has been deposited 
will have an excess of OH - ions and will appear pink due to the 
phenolphthalein. 

From this experiment and from the principles noted in Ex¬ 
ercise 78, draw conclusions as to the efficacy of zinc and tin in 
preventing the rusting of iron. 


132 


GENERAL CHEMISTRY 


EXERCISE 96 

PREPARATION OF DOUBLE SALTS 

A double salt may be formed by mixing solutions of the 
simple salts composing it in the proportions required by the 
formula of the double salt. Recall the double salts prepared 
in this manner in Exercises 77 (c), 87 (c) and 92. When crystals 
of the double salt are dissolved all the ions are given which 
would be formed by the component simple salts. This is not 
true of a complex salt such as K 4 Fe(CN) 6 . What ions does 
potassium ferrocyanide give? How should its formula be 
written if it were a double salt? 

(a) FeSOr (NH 4 ) 2 S0 4 6H 2 0. — Calculate the weight of iron 
which will dissolve in 50 cc. of dilute sulfuric acid (6N). Use 
about a gram in excess of this weight of clean iron wire, place 
it in 50 cc. of dilute sulfuric acid and allow the reaction to pro¬ 
ceed until no more hydrogen is evolved. Remove any iron 
which has not dissolved, filtering the solution if necessary. 

Neutralize another 50 cc. of dilute sulfuric acid with a con¬ 
centrated solution of ammonia and mix this solution with the 
solution of ferrous sulfate you have just prepared. Heat the 
mixture to boiling, evaporate it to about 75 cc., and set it aside 
to crystallize. Calculate the weight of the double salt crystals 
which could be formed from the amounts of reagents used. 

When the solution has cooled and a sufficient crop of crystals 
has appeared, filter out the crystals and reserve the filtrate for 
(b). Dry the crystals by exposing them to air on filter papers 
and weigh them. Measure the volume of the filtrate and cal¬ 
culate the weight of the double salt which it must still contain 
at 20° if the solubility of the salt is 22 parts per 100 parts 
of water at this temperature. Why will this calculation be 
slightly inaccurate even if the temperature of the filtrate is 
exactly 20° ? What weight of the crystals should you have 
obtained theoretically? What is your percentage yield? 

Are the crystals of this double salt efflorescent? Is the salt 
oxidized by air to a ferric compound? Test the crystals for 
ferrous and ferric iron, for the ammonium and sulfate ions, and 
for water of hydration. Write equations for all the chemical 
reactions which you have employed in this experiment? 


LABORATORY EXERCISES 


133 


(6) Fe 2 (S0 4 ) 3 - (NH 4 ) 2 S0 4 24H 2 0. — To the filtrate from (a) 
add, with care, 5 cc. of concentrated sulfuric acid. Heat the 
solution to boiling and add concentrated nitric acid, a little at 
a time, with stirring, until a drop of the solution in 2 cc. of 
water fails to give a test for the ferrous ion. 

Evaporate the solution to about one-half its volume, place a 
string across the beaker so that it hangs in the solution, and 
set the solution in your desk to crystallize. 

At the next laboratory period examine and describe the 
crystals. To what class of substances does this double salt 
belong? Write the formulas and a brief comparison of the two 
other substances of this type which you have prepared. 


EXERCISE 97 

COBALT 

Textbook: 728-730 

(a) Cobaltous Hydroxide. — Add a solution of sodium hy¬ 
droxide to a solution of a cobaltous salt. Describe the precipi¬ 
tate which is formed. Is it soluble in excess sodium hydroxide? 
Repeat the experiment with ammonium hydroxide in place of 
sodium hydroxide (?). Why does the precipitate dissolve in 
excess ammonium hydroxide? Note the color of the solution. 
Equation. 

(b) Cobaltous Sulfide. — Treat 10 cc. of a solution of a cobalt 
salt with hydrogen sulfide (?). Now add an excess of ammo¬ 
nium hydroxide to the solution and again treat it with hydrogen 
sulfide (?). Equation. To what group of sulfides does cobalt 
sulfide belong? Refer to Exercise 49 (6). Filter out the pre¬ 
cipitate, remove a small portion and test its solubility in dilute 
hydrochloric acid (1.2 normal) prepared by diluting 1 part of 
the concentrated acid with 9 parts of water (?). 

Prepare a borax bead as in Exercise 59 (6), take up a particle 
of the cobalt sulfide on this bead and heat it in the oxidizing 
flame of the Bunsen burner. The cobalt sulfide will be oxidized 
in the flame to cobalt oxide which will react with borax to pro¬ 
duce a mixture of the metaborates of cobalt and sodium. Note 
the color and write the equation. 

Transfer the remainder of the cobalt sulfide to a porcelain 


134 


GENERAL CHEMISTRY 


evaporating dish and dissolve it by boiling it with the minimum 
quantity of a mixture of 3 parts dilute hydrochloric acid with 
1 part dilute nitric acid (dilute aqua regia). Equation. Evap¬ 
orate this solution almost to dryness, take up the residue in 
water and filter it. This solution of cobaltous chloride should 
be reserved for use in (c). 

(c) Sympathetic Ink. — Dip a splint of wood or a glass rod 
in a dilute solution of cobalt chloride and write upon a piece 
of white paper. Permit this to dry and note whether or not the 
writing is visible. Warm the piece of paper over a flame being 
careful to avoid scorching the paper (?). Now breathe upon the 
paper (?). Explain the changes observed. To 1 cc. of a solution 
of cobalt nitrate add concentrated hydrochloric acid until a 
change in color is observed. Suggest the cause of this change. 

Many other sympathetic or secret inks can be devised. For 
example, writing with a solution of a lead salt can be developed 
by a blotter or filter paper moistened with a solution of hydro¬ 
gen sulfide; writing with potassium ferrocyanide can be de¬ 
veloped by means of a solution of a ferric salt. Suggest other 
possible combinations. 


EXERCISE 98 

NICKEL 

Textbook: 730-732 

(a) Nickelous Hydroxide. — To a solution of a nickelous 
salt add a solution of sodium hydroxide (?). Is the precipitate 
soluble in excess? Repeat the experiment substituting ammo¬ 
nium hydroxide as the precipitant (?). What is formed when 
excess ammonium hydroxide is added? The color resembles 
that of what other ammonia complex? Write the formulas for 
all the metal-ammonia ions which you have encountered in 
this course and note any characteristic colors. 

(; b ) Nickelous Sulfide. — Treat a solution of a nickel salt 
with hydrogen sulfide (?). Now make the solution alkaline wth 
ammonium hydroxide and again treat with hydrogen sulfidgM). 
Test the solubility of the precipitate in 1.2 normal hydrochloric 
acid and in dilute aqua regia (?). Compare the results with 
those obtained in Exercise 97 (b). Make a borax-bead test with 
nickel sulfide and compare it with the bead test for cobalt. 


LABORATORY EXERCISES 


135 


(c) Separation of Nickel and Cobalt. — Test separately 
small portions of solutions of nickel and cobalt salts with a so¬ 
lution of dimethyl-glyoxime. Describe the results and suggest 
a method for the separation and identification of nickel and 
cobalt. 


EXERCISE 99 

UNKNOWN NO. 5 —“GROUP UNKNOWN” 

The qualitative Groups I to IV are to be reported if pres¬ 
ent. No analysis of the precipitate is required. 


Group 


II 


III 


IV 






Procedure 


yc , 

Dil ute- - 4 -evrt7f~the~unkno WCi soluti on 
wi th 10 cc. of - water am i add 4 cc. of 
dilute HC1. Filter; test small portion 
oFlHtrate with HC1 to determine if 
precipitation was complete. 


Dilute filtrate to 100 cc. and saturate 
with H 2 S. Heat nearly to boiling and 
filter. 


To the filtrate from II add NH4OH to 
alkalinity (litmus test) and saturate 
with H 2 S. Filter. 


Acidify the filtrate from III with dil. 
HC1 and evaporate the solution to 25 
cc. in ah open beaker. Filter if a pre¬ 
cipitate (probably S) appears. Make 
the solution alkaline with NH 4 OH, 
heat to boiling, and add (NHthCCh 
solution. Allow to stand. 


The filtrate from IV would contain 
Na, K, NH 4 , and Mg as ions. There 
is no group reagent for Group V. It is 
not to be reported in this unknown. 


Precipitate 
May Contain 


AgCI, ^^ 3 ) 
Hg 2 Cl 2 ,*^ "d 1 

P bGhr- 




PbS, HgS, As 2 S 3 , 
"31^3, Bi 2 3 3 , 

Bnft , 

CuS, CdS. 


Al(OH) 3 , G**OH) 3 , 
FeS, MnSr 
ZnS, NiS 


CaC0 3 , 
SrC0 3 , 
BaC0 3 . 




















136 


GENERAL CHEMISTRY 


EXERCISE 100 

UNKNOWN NO. 6 — “GENERAL UNKNOWN” 

Not more than three cations (in all groups ) 

Procedure: The group separations are to be carried out as 
directed for Unknown No. 5. The group precipitates and the 
final filtrate (containing Group V) are to be analyzed as follows: 
Test for NH 4 salts using a portion of the original unknown. 


Group I 

Proceed as with Unknown No. 4 — Exercise 90 (d, e ). 

Group II — Separation of Tin and Copper Sub-groups 

(1) Wash the precipitate with water, transfer to a dish or 
beaker, add 10 cc. of (NH 4 ) 2 S X solution, stir and warm (do not 
boil) for 10 minutes. Add 10 cc. of water and filter. Residue 
is Copper Group, filtrate contains Tin Group. 

Filtrate — Analysis of Tin Group 

(2) Dilute and acidify with dil. HC1; boil, filter and wash 
precipitate. Discard filtrate, transfer precipitate to dish and 
heat it 10 minutes with 10 cc. of cone. HC1. Dilute, filter and 
wash the precipitate. 

Precipitate — As 2 S 5 and S 

(3) Add 10 cc. of dil. HN0 3 and warm. Dilute, filter, and 
evaporate the filtrate to 5 cc. Add magnesia mixture and cone. 
NH 4 OH until strongly alkaline. White precipitate on standing 
is MgNH 4 As0 4 . 

Filtrate — SbCl 3 and SnCL* 

(4) Dilute to exactly 50 cc., heat to boiling and saturate hot 
with H 2 S for 5 minutes. Filter and wash the precipitate. 

Precipitate — Sb 2 S 3 (orange) 

(5) Dissolve in cone. HC1, add aluminum wire. Black pre¬ 
cipitate insoluble in HC1 is Sb. 



LABORATORY EXERCISES 


137 


Filtrate — SnCl 4 

(6) Evaporate to 30 cc., add one piece A1 wire and heat until 
all action ceases. Filter into a solution of HgCl 2 . White or 
gray ppt. shows Sn. 


Residue (from 1) — Analysis of Copper Group 

(7) Add 5 cc. of dil. HN0 3 to 10 cc. of water, add this mix¬ 
ture to the residue in a beaker and boil for several minutes. 
Filter, wash ppt. 

Residue (from 7) — HgS (black) and S 

(8) Boil with 3 or 4 cc. of aqua regia until N0 2 is no longer 
evolved, dilute with 10 cc. of water and filter. Add SnCl 2 to 
filtrate (?). 

Filtrate — Nitrates of Pb, Bi, Cu, Cd 

(9) Add 3 cc. of cone. H 2 S0 4 and evaporate in a beaker until 
dense white fumes of S0 3 appear. Allow to cool and pour into 
15 cc. of water — rinse out first beaker with this solution. Stir 
and allow to stand; filter. 

Precipitate — PbS0 4 (white) 

(10) Dissolve ppt. in hot ammonium acetate solution and add 
K 2 Cr 2 0 7 . A yellow ppt. insoluble in HC 2 H 3 0 2 is PbCr0 4 . 

Filtrates — Sulfates of Bi, Cu, Cd 

(11) Add NH 4 OH, stirring well, until the solution is al¬ 
kaline — then add 2 cc. NH 4 OH in excess. Filter. Precipitate 
is Bi(OH) 3 . Add Na 2 Sn0 2 — black residue is Bi. 

Filtrate — Cu and Cd as ammonia-complex sulfates 

(12) If colorless, copper is absent — proceed with (14). If 
blue, save most of solution for (13) — acidify a small portion 
with H 2 S0 4 and add 2 drops of K 4 Fe(CN) 6 . Brown precipitate 
is Cu 2 Fe(CN) 6 . 

(13) Add an excess of dil. H 2 S0 4 and about 5 g. of iron filings, 
cover with watch glass and boil several minutes. Filter and 
treat filtrate by (14). 



138 


GENERAL CHEMISTRY 


Filtrate — Cd or Cd(NH 3 ) 4 as sulfate 

(14) Acidify if necessary and saturate with H 2 S. Yellow 

precipitate is CdS. Explain how Cu but not Cd was removed 
in (13) and why Fe ions do not interfere with the test for Cd 
in (14). ____ 

Group III — Separation of Fe and A1 Sub-groups 
Precipitate — Al(OH) 3 , Cr(OH) 3 , ZnS, MnS, FeS, CoS, NiS 

(15) Transfer to beaker and add dil. HC1. If black residue 
(NiS, CoS) remains add a few drops of cone. HN0 3 and boil. 
Add 10 cc. water and filter out S. Evaporate filtrate to 10 cc. 
Add NaOH in excess and then Na 2 0 2 slowly with stirring. Boil, 
dilute, filter and wash residue. Analyze filtrate by (16, 17); 
residue by (18-20). 

Aluminum Group 

Filtrate — NaA10 2 , Na 2 Cr0 4 , Na 2 Zn0 2 

(16) Add HN0 3 to acidify, then NH 4 OH to alkalinity, stirring 
well. Heat to boiling and filter. White precipitate is Al(OH) 3 . 

Filtrate — Na 2 Cr0 4 (yellow), Zn(NH 3 ) 4 ions 

(17) Add HC 2 H 3 0 2 to acidity — Na 2 Cr 2 0 7 and Zn(C 2 H 3 0 2 ) 2 
are formed. Yellow or orange color indicates Cr — confirm by 
test in Exercise 91 (6). To the larger portion of the solution 
add 2 drops K 4 Fe(CN) 6 and warm. A white precipitate is 
Zn 2 Fe(CN) 6 . 

Iron Group 

Residue (from 15) — Mn0 2 , Fe(OH) 3 , Ni(OH) 2 , Co(OH) 3 

(18) Transfer to a porcelain dish, add 15 cc. cone. HN0 3 and 
heat to boiling. If a black residue remains add NaN0 2 solution 
until residue dissolves. Boil until N0 2 is no longer evolved. 
Add gradually, with stirring, 1 g. solid KC10 3 and boil for 
3 minutes. A brown or black residue is Mn0 2 . Filter through 
a plug of asbestos wool and wash. Transfer precipitate to a 
beaker, add 1 g. Pb0 2 and 10 cc. dil HN0 3 . Boil and allow 
solids to settle. Pink color indicates permanganate. 

Filtrate — Nitrates of Fe, Co, Ni 

(19) Add NH 4 OH in excess, filter. Precipitate is Fe(OH) 3 . 
Dissolve it in dil. HC1, add NH 4 SCN. Red solution is Fe(SCN) 3 . 



LABORATORY EXERCISES 


139 


Filtrate — Ni, Co as ammonia-complex ions 

(20) Acidify with HC2H3O2, divide into two portions: 

(A) Add dimethyl glyoxime. Red precipitate indicates 
Ni; brown solution Co. 

(B) Evaporate second portion to 2 cc., and add 2 cc. 
NH 4 SCN and 2 cc. amyl alcohol-ether mixture and 
shake. Blue-green layer indicates Co. 


Group IV — CaC0 3 , SrC0 3 , BaC0 3 

(21) Dissolve the precipitate by repeatedly pouring 10 cc. 
of hot HC2H3O2 through the filter and analyze the solution as 
in Exercise 81 (c, d ). 


Group V — Mg, K, Na, NH 4 Salts 
Filtrate from Group IV 

(22) Add 5 cc. cone. HN0 3 and evaporate to dryness in a 
small beaker. Moisten residue with cone. HN0 3 and ignite to 
decompose NH 4 NO 3 . Cool, add 10 cc. water, a few drops HC1 
and filter if not clear. Test one-third of filtrate by (23) and two- 
thirds by (24). 

(23) Add cone. NH 4 OH through a filter until solution is 
alkaline. If a precipitate forms dissolve it by adding 3 cc. 
NH 4 CI solution through the filter. Then add 2 cc. Na 2 HP0 4 
through the filter, shake the mixture vigorously and allow to 
stand. A white ppt. is MgNH 4 P0 4 . 

(24) Make flame tests and precipitation tests for Na and K 
as in Exercise 73. 


If any difficulties are met consult a manual of Qualitative 
Analysis or an instructor. 

Additional unknowns consisting of alloys, everyday sub¬ 
stances, etc., may be given for further practice in qualitative 
analysis. 











PART III 


CHEMICAL ARITHMETIC 

TYPES OF PROBLEMS 

The first thing that is necessary in the working of any chemi¬ 
cal problem is that the student should have clearly in mind the 
full significance of a chemical formula. From the text it is to 
be seen that the formula for potassium chlorate, KC10 3 , stands 
for one gram molecule of KCIO 3 which is made up of one gram 
atom of potassium, one gram atom of chlorine and three gram 
atoms of oxygen. Before a quantitative problem can be worked 
these amounts must be converted to actual weight units, that 
is, we must think of KC10 3 as standing for a compound that 
is made up of 39.1 grams of potassium, 35.46 grams of chlorine 
and 3 X 16 grams or 48 grams of oxygen. When this conversion 
has been made, the material used in the experiment can be 
followed quantitatively from change to change through any 
series of balanced equations. All chemical problems fall under 
one of the following classes: 

A. Problems Involving Weight Relationships Alone: 

I. Percentage Composition. 

II. Chemical Formula from Percentage Composition. 

III. Amount of One Substance Required to Combine with 
or Displace a given Amount of Another Substance. 

B. Problems Involving Weight and Volume: 

IV. Specific Gravity of Solids and Liquids. 

V. The Correction of Gas Volumes for Temperature, 
Pressure and Water Vapor. 

VI. Gram-Molecular Volume. 

VII. Normal Solutions. 

Before starting to work a problem analyze it carefully and 
classify it under one or more of the above types. It may be 
that the problem may involve the use of two or more of the 
above types in its solution. 


141 


142 


GENERAL CHEMISTRY 


I. Percentage Composition from Formula. 

A formula represents a molecule or a gram molecule of a sub¬ 
stance. It indicates the kind and number of atoms of which the 
molecule is composed. If the constituent atoms are all alike, 
the substance is an element — e.g., 0 2 , Cl 2 . If the atoms present 
in a molecule are those of different elements, the substance is a 
compound — e.g., H 2 0, KC10 3 . Since to each element a definite 
relative (atomic) weight may be assigned, the percentage by 
weight of any element in a compound may be calculated from 
the formula of the compound. 


Model Example I 

To calculate the per cent of oxygen in potassium chlorate. 

Method of Solution. — The formula of potassium chlorate, 
KCIO 3 , shows that it contains an atomic weight of potassium 
(39), one of chlorine (33.5), and three atomic weights of oxygen 
(3 X 16). The sum of these numbers is 122.5 and is called 
the molecular weight of potassium chlorate. To express this 
number in actual weight units one may hence say that a gram 
molecule of potassium weighs (or is) 122.5 g. Of this weight 
48 parts are oxygen. The percentage of oxygen in the compound 
is, therefore, the ratio to 100 which is equal to the ratio of 48 
to 122.5. 

That is, x : 100 = 48 : 122.5 and hence x = 39.18 per cent. 

1 . Compare the percentages of iron in ferrous oxide, ferric 
oxide, and magnetic oxide of iron. Which of the three is the 
most valuable iron ore? The formulas are: FeO, Fe 2 0 3 , Fe 3 0 4 . 

2 . (a) How many different kinds of atoms are there in each 
of the following substances: C 2 H 5 OH, (NH 4 ) 2 S0 4 , Ci 2 H 22 0n, 
Fe 2 (S0 4 ) 3 ? 

( 6 ) How many atoms are there in each of these molecules? 

(c) Calculate the molecular weight of each substance. 

3. Calculate the per cent of oxygen in each substance in 2 (a). 

4. Calculate the percentage of water in crystallized “ washing 
soda,” Na 2 CO 3 -10H 2 O. 

5. How many grams of sodium are present in 150 grams 
of sodium bicarbonate, NaHCCb? 

6 . What weight of phosphorus is present in a kilogram of 
calcium phosphate, Ca 3 (P0 4 ) 2 ? 


CHEMICAL ARITHMETIC 


143 


II. Derivation of Formula from Percentage or Weight Com¬ 
position. 

When the composition of a substance is given in percentages 
or in actual weights, the ratio of the atoms can be determined 
and the simplest formula can thus be derived. If the molecular 
weight of the compound is known, the actual formula can be 
found for the compound. 

Model Example II 

A compound contains 52.18 per cent carbon; 13.04 per cent 
hydrogen; and 34.78 per cent oxygen. What is its simplest 
formula? 

Method of Solution. — The ratio of each kind of atom pres¬ 
ent is obtained by dividing the amounts of the elements by 
their respective atomic weights — thus: 

CO IQ 

c = = 4.348; divided by 2.174 = 2 

H = — -° 4 = 13.04; divided by 2.174 = 6 

Q4 7R 

O = ^ - = 2.174; divided by 2.174 = 1 

The first quotients which represent the atomic ratios must 
now be convert^! to small whole numbers. This is accomplished 
by djpiding them by their highest common factor. It is ap¬ 
parent that the quotients obtained above are all multiples of 
2.174 and hence the ratio of the atoms in the compound is 

C : H : O = 2 : 6 : 1 and the simplest formula is C 2 H 6 0. If 

this is the correct formula the molecular weight must be 
2 X 12 + 6 X 1 + 16 = 46. If, however, the molecular weight 
is some multiple of 46 ( e.g ., 92) the formula must be multiplied 
accordingly {e.g., C 4 H 12 O 2 for a mol. wt. of 92). 

7. A compound having a molecular weight of 30 contains 80 
per cent carbon and 20 per cent hydrogen. What is its formula? 

8 . Ten grams of iron when burned in oxygen form 13.81 g. 
of an oxide. Calculate the formula of the oxide. Assuming 
this to be the correct formula, calculate the molecular weight 
of the compound. 

9. A compound by analysis gave 26.9 per cent sodium, 





144 


GENERAL CHEMISTRY 


16.58 per cent nitrogen, and 56.52 per cent oxygen. Calculate 
its simplest formula. 

10. A compound having a molecular weight of 106.5 is found 
to contain 21.6 per cent sodium, 33.3 per cent chlorine, and the 
remainder is oxygen. What is its simplest formula? Is this 
also its actual formula? 

11. Derive the empirical formulas for the substances which 
have the following percentage compositions: 

(a) K = 26.59% ( b ) K = 40.26% 

Cr = 35.39 Cr = 26.77 

O = 38.02 O = 32.95 

12. Derive the empirical formula of the substance which 
gave the analysis*; Ca = 18.28 per cent, Cl = 32.36 per cent, 
H 2 0 = 49.36 per cent. 

III. Simple Weight Calculation Based upon a Chemical Re¬ 
action. 

In this type of problem the correct chemical equation must 
always be written first, and the calculation then based upon 
the weights of the substances involved in the reaction. 


Model Example III 


What weight of oxygen may be obtained by the action of 
water upon 25 g. of sodium peroxide? 

Method of Solution. — The equation, ^ 

2Na 2 02 + 2H 2 0 = 4NaOH + " 0 2 \ 

2(2 X 23 '+ 2 X 16) g. 2^06g! 


shows that two gram molecules, 156 g., of sodium peroxide.are 
required to produce one gram molecule, 32 g., of oxygen. Un¬ 
less the equation had been written it might have been wrongly 
assumed that sodium peroxide would liberate all of its oxygen 
instead of only half of it. Since the equation shows that 156 g. 
of Na 2 0 2 will give 32 g. of 0 2 , it must follow that 25 g. of 
Na 2 0 2 will give a proportionate amount — i.e., 


156 g. 32 g. ,t. 32 X 25 

vte— = -; and rX = ■ — 

25 g. x g. ’ 156 


5.12 g. 0 2 


Care must be taken to have the proportion correctly expressed. 
It should also be noted that only two kinds of quantities can 




CHEMICAL ARITHMETIC 


145 


ever appear in any single proportion. In one ratio of the 
above proportion we have grams of sodium peroxide; in the 
other ratio we have grams of oxygen. The weight of water 
used in the reaction or the weight of sodium hydroxide pro¬ 
duced could be calculated by constructing separate proportions 
for each. 

13. What weight of mercuric oxide must be decomposed in 
order to obtain 50 g. of oxygen? 

14. What weight of oxygen may be obtained from 100 g. of 
potassium chlorate? 

15. What weight of potassium chloride is formed when 20 g. 
of oxygen are liberated from potassium chlorate? 

16. What weight of oxygen is needed to burn completely 40 g. 
of phosphorus? 

17. How many grams of oxygen are required to burn 160 g. 
of sulfur? How many gram atoms of oxygen is this? What 
weight ofWilfur dioxide is formed? This weight of sulfur 
dioxide will^ombine with what weight of water? 

18. What weight of magnetic oxide of iron is formed when 
56 g. of iron are burned in oxygen? What weight of oxygen 
combines with 56 g. of iron? 

19. Homuch zinc sulfate is formed when a gram of hydro¬ 
gen is liberated by the action of sulfuric acid on zinc? 

20. Calculate the weight of hydrogen obtained when 10 g. 
of water react jyith sodium. What weight of hydrogen is ob¬ 
tained from the same weight of water reacting (as steam) 
with'iron? 

(21. Hydrogen is passed over hot magnetic oxide of iron until 
therlatter has lost 3.2 g. What weight of metallic iron has been 
produced? 

(22, What weight of water is formed by reducing 80 g. of 
copper oxide with hydrogen? What weight of copper is left? 

IV. Specific Gravity of Solids and Liq^ds. 

The density (mass per unit volume) of solids and liquids is 
usually compared with that ofj^feter as the standard. The 
relative density compared wi^Rvater is called the specific 
gravity. It is customary to make this comparison at the tem¬ 
perature at which water has its maximum density (4° C.) and 
at standard atmospheric pressure. When a body is heavier than 


146 


GENERAL CHEMISTRY 


water, the specific gravity may be obtained by dividing its 
weight in air by its apparent loss of weight when weighed in 
water. If the substance is soluble in water, it may be weighed 
in a liquid in which it is insoluble and of which the specific 
gravity is known. 


Model Example IV 

A piece of granite weighs 2.30 g. in air and 1.52 g. in water. 
Calculate the specific gravity of the granite. 

Method of Solution. — The sample of granite displaces a vol¬ 
ume of water equal to the volume of the sample of granite. The 
apparent loss in weight of the sample of granite is equal to the 
weight of the water displaced. Since 1 cc. of water weighs 
1 g., the loss in weight of the granite is numerically equal to 
the volume of the water displaced which is, in turn, equal to the 
volume of the sample of granite. In the problem above, the loss 
in weight is 0.78 g., hence the volume of the granite is‘ 0.78 cc. 
and its specific gravity is 2.30/0.78 = 2.948. 

23. What volume is occupied by 100 g. of HN0 3 if its specific 
gravity is 1.42? 

24. A solution of hydrochloric acid has a specific gravity 
of 1.2 and contains 40 per cent by weight of HC1. Calculate 
the volume of this acid solution which will contain 50 g. of 

HC1. 

25. A piece of glass weighs 25.86 g. in air, 15.26 g. in water, 
and 9.20 g. in sulfuric acid. Calculate the specific gravity of 
the glass and of the sulfuric acid. 

26. A bottle weighs 50.50 g. When filled with water the 
bottle weighs 258.58 g., and when filled with oil it weighs 
220.20 g. Calculate the specific gravity of the oil. 

27. A graduate cylinder is filled with water to the 75-cc. 
mark. When 25 g. of rhombic sulfur are dropped into this, 
the water level rises to Jhe 87.2-cc. mark. Calculate the specific 
gravity of rhombic sulrar. 

28. What volume of sulfuric acid of specific gravity 1.84 and 
containing 94 per cent by \mght of H 2 S0 4 will be required to 
neutralize completely 50 cc. of sodium hydroxide solution which 
has a specific gravity of 1.35 aim contains 31.5 per cent by weight 
of NaOH? 


CHEMICAL ARITHMETIC 


147 


V. Application of the Laws of Boyle and Charles — Calculation 
of Volume Change Caused by Change in Temperature and 
Pressure. 

The volume occupied by a definite quantity of gas at a cer¬ 
tain temperature and pressure may be calculated if its volume 
at any other temperature and pressure is given. It is only 
necessary to remember that the volume of a gas varies directly 
with the absolute temperature and inversely with the pressure 
to which it is subjected. 

The volume occupied by a gas has no significance unless the 
temperature and pressure at which it is measured are also 
stated. For convenience a definite temperature and pressure 
have been agreed upon at which to compare gases. These are 
0° Centigrade and 760 mm. pressure and are called normal or 
standard conditions. Unless others are specified, these condi¬ 
tions are understood when the volume of a gas is mentioned. 

Gases are ordinarily measured over dry mercury because that 
liquid has a very low vapor pressure at ordinary temperatures. 
If a gas is measured over water, part of its apparent volume 
and pressure is due to w^ter vapor, for which correction must 
be made in calculating the true volume of the gas. 


Model Example V 


If a quantity of oxygen measures 250 cc. at 18° C. and 750 
mm. pressure, what volume will it occupy at 25° C. and 780 mm. 
pressure? 

Method of Solution. — First convert the temperatures to the 
Absolute scale by algebraically adding 273°. We can now apply 
the two gas laws. In the change indicated in the problem the 
temperature is increased and will increase the volume in the 


273 + 25 

ratio of-(For an increase, the larger number is 

273 + 18 

placed on top.) The pressure is also increased but this will 

750 


tend to decrease the volume in inverse ratio — ^.e., 


780 


The 


two effects can be combined in tl|e form: 

250 cc X — X — = ? cc. at the new conditions, 25° C., 

291 780 


780 mm. 



148 


GENERAL CHEMISTRY 


If the gas were measured over water at 18° C. and 750 mm. 
and its dry volume at 25° C. and 780 mm. is desired, it is 
necessary to subtract the vapor pressure of water at 18° (refer 
to table) from the observed pressure of 750 mm. and to use the 
remainder in place of 750 in the above calculation. 

29. If a volume of gas measures 1 1. at 0° C., what will be 
its volume at 100° C. if the pressure remains constant? 

30. If the volume of a gas measures 400 cc. at 750 mm. 
pressure, calculate its volume if the pressure is changed to 
500 mm., assuming constant temperature. 

31. A volume of air measuring 200 cc. at 20° C. and 700 mm. 
pressure will become what volume at — 20° C. and 800 mm. 
pressure? 

32. Calculate the volume at standard conditions occupied by: 

(a) 500 cc. of a gas measured at 22° C. and 640 mm. pressure. 

(b) 310 cc. of a gas measured at - 15° C. and 900 mm. pres¬ 
sure. 

(c) 25 1. of gas measured at 100° C. and 100 mm. pressure. 

33. Twenty liters of hydrogen at' standard conditions will 
occupy what volume over water at 26° C. and 765 mm. pressure? 

34. If an automobile tire has a volume of 1.2 cu. ft. when 
inflated, how many cu. ft. of air, measured at standard condi¬ 
tions, must be pumped into it in order to give a pressure of 3 
atmospheres at 25° C.? An atmosphere of pressure is 760 mm. 
of mercury or 15 lb. per square inch. 

35. What volume of hydrogen at standard conditions will be 
liberated from a gram molecule of sulfuric acid by zinc? What 
volume will this hydrogen occupy at 20° C. and 870 mm. 
pressure? 

36. If 430 cc. of a gas at 20° C. and 780 mm. pressure weigh 
1.54 g., what is the weight of a liter of the gas at standard 
conditions? 

37. Oxygen was collected over water when the barometric 
pressure was 735 mm. and the temperature was 22° C. What 
would be the true volume erf 200 cc. of this gas at standard 
conditions? 

38. What weight of zinc would be required to displace enough 
hydrogen from an acid to inflate a balloon having a capacity of 
500 liters at 20° C. and 800 mm. pressure? 


CHEMICAL ARITHMETIC 149 

VI. Problems Involving the Gram-Molecular Volume of Gases. 

It follows from Avogadro’s Law that gram molecules of dif¬ 
ferent gases will occupy equal volumes. The volume of a gram 
molecule of any gas at standard conditions is 22.4 liters. Hence 
the molecular formula of any gas has a volume significance as 
well as one of definite weight — see Type III. For example, 
the formula 0 2 may be read 32 g. of oxygen or 22.4 1. of oxygen 
(standard conditions); the formula C0 2 represents 44 g. and 
also 22.4 1. of carbon dioxide gas. 

Model Example VI 

What volume of oxygen at standard conditions may be ob¬ 
tained from 50 g. of potassium chlorate? 

Method of Solution. — This problem, like Model III, is based 
upon a chemical reaction, so a balanced equation should first be 
written: 

2KC1Q 3 = 2KC1 + 30 2 

2 X 122.5 g. 3 X 22.4 1. 

This shows that 245 g. of potassium chlorate will liberate 

67.2 1. of oxygen. A proportionate volume of oxygen may be 

obtained from 50 g. of the chlorate — i.e., 

245 67.2 10Q1 , 

-3TT- -->; x = 13.81. of oxygen. 

ou x 

( Caution: It must be noted that the gram-molecular volume 
(G.M.V.) idea, like Avogadro’s Law, applies only to sub¬ 
stances in the gaseous state. Also, care must be taken to em¬ 
ploy the correct molecular formula for the gas. Thus, H 2 , and 
not H, represents 22.4 1. of hydrogen.) 

39. If 90 g. of water are decomposed by electrolysis, what 
volume of hydrogen is obtained? What volume and weight of 
oxygen? If the same weight of water reacts with sodium, what 
volume of hydrogen is obtained? 

40. What weight of iron will be produced in the reduction 
of its magnetic oxide by means of 24 1. of hydrogen? 

41. What volume of ozone may be obtained from 75 1. of 
oxygen if the conversion is 8 per cent efficient? What will this 
volume of ozone weigh? 

42. What is the numerical value of the G.M.V. at 25° C. 
and 750 mm. pressure? 




150 


GENERAL CHEMISTRY 


43. How many grams of oxygen are needed to burn com¬ 
pletely 60 g. of carbon? How many gram atoms of oxygen is 
this? How many gram molecules? What volume does the 
product occupy at standard conditions? 

44. What is the weight of a liter of a gas having the formula 
C 2 H 6 ? What is its density compared with air? 

45. If 400 cc. of a gas weigh 0.520 g., what is the molecular 
weight of the gas? 

46. If an electric spark is passed through a mixture of 524 cc. 
of hydrogen and 275 cc. of oxygen, will any gas remain after the 
tujjg^is cooled? If so, what gas and how much? 

'47, What volume of hydrogen must be burned in chlorine in 
order to make enough hydrochloric acid to neutralize 60 g. of 
sodium hydroxide? 

487" 'Compare the volumes of chlorine produced when hydro¬ 
chloric acid reacts with 50 g. of manganese dioxide and when it 
reacts with 50 g. of potassium permanganate. 

49. What volume of chlorine may theoretically be liberated 
from a kilogram of bleaching powder by the action of sulfuric 
acid? 

50. What volume of chlorine, reacting with a dilute solution 
of KOH, is needed to prepare 12 g. of potassium hypochlorite? 
What volume of chlorine, reacting with a concentrated solution 
of KOH, is needed to prepare 12 g. of potassium chlorate? In 
each case what per cent of the chlorine used appears in the 
product named? 

51. What volume of bromine vapor, measured at 25° C. and 
780 mm. pressure, may be formed from 200 cc. of liquid bromine? 
The density of liquid bromine is 3.18. 

52. What volume of hydrogen bromide is obtained when 10 g. 
of phosphorus are combined with bromine and the product 
hydrolyzed with water? 

53. A 10 per cent solution of potassium iodide has a density 
1.08. What volume of chlorine is needed to displace the iodine 
in a liter of this solution? 

X 54. What volume of hydrogen fluoride, H 2 F 2 , is needed to 
react with 20 g. of CaSi0 3 ? What weight of calcium fluoride is 
required to prepare this quantity of hydrogen fluoride? 

55. Calculate the weight of ferrous sulfide needed to prepare 
enough H 2 S to fill a gasometer having a capacity of 2 cu. m. at 
25° C. and 750 mm. pressure. 



% 

CHEMICAL ARITHMETIC 151 

* 56. What weight of antimony sulfide is formed when 2 1. of 
H 2 S are passed into a solution of antimony chloride? 

57. If a solution of K 2 Cr 2 07 is acidified with HC1 and 5 1. of 
H 2 S gas are used, what weight of K 2 Cr 2 0 7 will be reduced? 

58. What weight of copper is needed to prepare 10 1. of sulfur 
dioxide by the reduction of concentrated sulfuric acid? What 
weight of sodium sulfite may be prepared with this quantity of 
the gas? 

59. What volume of sulfur dioxide will, when dissolved in 
water, reduce 30 g. of potassium permanganate? 

60. Compare the volumes of sulfur dioxide obtainable from 
5 g. of Na 2 S0 3 and from 5 g. of Na 2 S 2 0 3 -5H 2 0 by the action of 
an acid. 

61. What volume of hydrogen will combine with 30 1. of 
nitrogen in the Haber synthesis of ammonia? What volume of 
ammonia will be produced by this combination? 

62. What volume of ammonia is liberated by the action of 
lime upon a kilogram of ammonium sulfate? By the hydrolysis 
of a kilogram of calcium cyanamide? 

VII. Equivalent Weights and Normal Solutions. 

The following problems deal with solutions of definite strength, 
expressed in terms of normality — that is, the number of gram 
equivalents of the solute in a liter of the solution. The nor¬ 
mality idea may involve: 

(a) The weight of a substance and the volume and strength 
of a solution. 

(5) The volumes and strengths of two solutions. 

(c) The volume of a gas and the volume and strength of a 
solution. 

Model Example VII 

(a) What weight of pure H 2 S 04 is contained in 400 cc. of 
1.2 N sulfuric acid solution? 

(6) If 35 cc. of this solution just neutralize 35 cc. of a basic 
solution, what is the normality of the latter? 

(c) What volume of C0 2 gas will be liberated by the action 
of 200 cc. of this 1.2 N acid on marble? 

Method of Solution. — (a) A gram molecule of sulfuric acid 
(98 g.) contains 2 gram equivalents. In a liter of normal sul- 
mfic acid there will hence be 49 g. of H 2 S 04 . In a liter of 1.2 N 


152 


GENERAL CHEMISTRY 


sulfuric acid there will be 49 X 1.2 = 58.8 g. In 400 cc. of 1.2 N 
sulfuric acid there will be 58.8 X = 23.52 g. of H 2 S0 4 . 

(6) It is apparent that 1 cc. of the base is equivalent to 
35/25, or 1.4 cc. of the acid. That is, the base is 1.4 times as 
strong as the acid. The acid is 1.2 N, so that the base must be 

1.2 N X 1.4 = 1.68 N. 

It may be noted that the product of normality X volume of 
any solution = the product of normality X volume of an 
equivalent amount of any other solution. That is, NV = N'V'. 
If three factors in this equation are known, the fourth may be 
found. In the above problem the unknown is the normality 
(N') of one solution. Using the equation to solve this problem, 


1.2 X 35 
25 


1.2 X 35 = N' X 25; N' = 


= 1.68 N. 


(c) The chemical action involved may be expressed by the 
equation: 

CaC0 3 + H 2 S0 4 = CaS0 4 + H 2 0 + C0 2 . 

This shows that 2 1. of N (or 1 1. of 2 N) H 2 S0 4 will liberate a 
gram-molecule, 22.4 1. of C0 2 gas. It follows that 11. of N acid 
will liberate 11.2 1. of C0 2 ; that 1 1. of 1.2 acid will liberate 

11.2 X 1.2 = 13.44 1. of C0 2 ; and hence 200 cc. of 1.2 acid will 

liberate 13.44 X = 2.688 1. or 2,688 cc. of C0 2 . 

63. A gram molecule of A1 2 (S0 4 )3 will make how many liters 
of N/5 solution? What weight of the salt will be present in a 
liter of this solution? 

64. The specific gravity of a 20 per cent solution of hydro¬ 
chloric acid is 1.1. What is the normality of the solution? 

65. How many grams of the following substances will be 
present in a liter of normal solution of each: H 3 As0 4 , SnCl 4 , 
Fe(N0 3 ) 3 ? 

66. How many gram equivalents of Na 2 S0 4 are needed to 
prepare 5 1. of 2 N sodium sulfate solution? How many grams 
of sodium sulfate will this be? 

67. What volume of 1.5 N base will be needed to neutralize 

3 1. of 6 N acid? 

68. What weight of sodium hydroxide is needed to neutralize 

4 1. of 1.5 N sulfurous acid? What volume of 0.2 N solution will 
this weight of NaOH prepare? 







CHEMICAL ARITHMETIC 


153 


69. If 500 cc. of an acid liberate 1.1 1. of S0 2 from a sulfite, 
what is the normality of the acid? 

70. If a gram molecule of S0 2 is dissolved in enough water 
to make 10 1. of sulfurous acid solution, what is the normality 
of the solution? 

71. What weight of lead sulfide may be precipitated from a 
solution of lead nitrate by the action of 500 cc. of N/5 H 2 S 
solution? 

72. What volume of ammonia gas will be liberated from an 
ammonium salt by the action of 3 1. of 1.3 N alkali? 

73. What volume of ammonia gas is required to neutralize 
300 cc. of 2.5 N phosphoric acid? 

74. If the ammonia resulting from the action of a base on 
4 g. of ammonium nitrate is dissolved to make 500 cc. of am¬ 
monium hydroxide, what is the normality of the solution? 


MISCELLANEOUS EXERCISES 


Valence Exercises 


75. Indicate the valence of the metal and of the acid radical 
in each of the following: NaN0 3 , MgC0 3 , A1C1 3 , Ca 3 (P0 4 ) 2 , 
K 2 S0 4 , Bi 2 S 3 , AgC10 3 , Cr 2 (S0 4 ) 3 . Name each compound. 

76. Using the valences found in the preceding problem, write 
the formulas for: calcium nitrate, sodium sulfide, aluminum 
phosphate, chromium chlorate, bismuth chloride, silver sulfate, 
potassium carbonate, magnesium chlorate. 

77. Write formulas for the salts formed by the action of nitric, 
sulfurous and phosphoric acids on each of the following bases: 
KOH, FeO, Fe(OH) 3 , Ca(OH) 2 . Name the twelve salts. 

78. If mercury and tin form the oxides: HgO, Hg 2 0 and SnO, 
Sn0 2 , write the formulas for mercurous nitrate, mercuric chlo¬ 
rate, stannic chloride and stannous sulfate. 


79. What weight of copper must react with nitric acid to 
produce 30 1. of nitric oxide? 

801 What volume of nitrogen dioxide will be formed by the 
action of concentrated nitric acid on 12 g. of copper? 

81. What weight of ammonium nitrate can be obtained by 
neutralizing 3 1. of 1.4 N nitric acid with ammonia? What 
volume of nitrous oxide can be obtained from this weight of 
ammonium nitrate? 


154 


GENERAL CHEMISTRY 


82. What weight of nitric acid is present in a liter of its 
solution, specific gravity 1.42, containing 70 per cent HN0 3 ? 
What volume of ammonia will this weight of nitric acid neu¬ 
tralize? 

83. What weight of carbon is required to reduce enough bone 
ash, mixed with sand, to form 12 g. of phosphorus? 

84. What volume of phosphine can be prepared by the action 
of 400 cc. of 3 N KOH on phosphorus? 

85 i What volume of chlorine gas will combine with 6 g. of 
phosphorus to form the pentachloride? What volume of N HC1 
solution may be prepared by the hydrolysis of this product? 

86. What volume of N/2 H 2 S solution will be required to 
precipitate 8 g. of antimony sulfide from a solution of SbCl 3 ? 

87. What weight of Na 2 C0 3 and what volume of 3 N H 2 S 04 
must be mixed to produce 2 1. of carbon dioxide? 

88. Calculate the volume of N/4 calcium hydroxide that will 
react with 2 1. of carbon dioxide. 

89. What volume will 20 g. of C0 2 occupy at 25° C. and 775 
mm. pressure? 

90. What weight of magnesium is needed to reduce 6 g. of 
silica to magnesium silicide? 

91. Calculate the weight of a G.M.V. of air assuming it to 
be a mixture of: nitrogen, 78 per cent; oxygen, 21 per cent; 
argon (molecular weight 40), 1 per cent. 

92. If a gas has a density of 1.24 referred to air, calculate the 
weight of a liter of the gas and its molecular weight. 

93. If 654 cc. of hydrogen, measured at 23° C. and 745 mm. 
pressure, are liberated from 200 cc. of sulfuric acid solution by 
the action of magnesium, calculate the normality of the sulfuric 
acid. 

94. Calculate the percentage of water of hydration in borax. 
What weight of crystalline borax must be treated with an acid 
to prepare 5 g. of boric acid? 

95. If a certain ore contains 30 per cent cuprous oxide, what 
weight of the ore must be worked in order to obtain a ton of 
pure copper? 

96. What volume of sulfur dioxide is obtained when a kilo¬ 
gram of iron pyrites is roasted? What volume of oxygen is re¬ 
quired? What volume of air? What weight of carbon will be 
required to reduce the iron oxide which is produced from a 
kilogram of iron pyrites? 


CHEMICAL ARITHMETIC 


155 


97. The specific heat of platinum is 0.0324, and a gram atom 
of chlorine combines with 48.6 g. of platinum to form platinic 
chloride. Calculate ( a ) the atomic weight of platinum and 
(b) the formula of platinic chloride. 

98. What volume of hydrogen is obtained when 10 g. of 
magnalium alloy (30 Mg, 70 Al) dissolve in sulfuric acid? 

99. What weight of pure nitric acid is required to dissolve 
100 g. of coinage alloy (90 Ag, 10 Cu) assuming that nitric 
oxide is formed in the reactions? What volume of 5 N nitric 
acid solution would this be? 

100 . What weight of microcosmic salt must be heated to pre¬ 
pare 5 g. of sodium metaphosphate? What volume of ammonia 
is ljjber.ated at the same time? 

101. What volume of arsine will be formed if a gram of AS 2 O 3 
is dissolved in HC1 and added to a hydrogen generator? 

, 102 . What volume of carbon dioxide can be obtained (a) by 
heating 42 g. of sodium bicarbonate? ( b ) by treating the same 
weight of the salt with an acid? 

103. Calculate the weight of caustic soda and the volumes of 
chlorine and hydrogen which may be obtained by the electroly¬ 
sis of a solution containing a kilogram of common salt. 

104. Twenty cc. of a sodium carbonate solution required 
50 cc. of N/1’0 hydrochloric acid for neutralization to the phenol- 
phthalein end-point. What volume of this acid would have 
been required if methyl orange had been used as the indica¬ 
tor? What weight of Na 2 C0 3 is contained in a liter of this 
solution? 

105. A mixture of Na 2 C0 3 and NaHC0 3 was dissolved in 
water and titrated with normal hydrochloric acid. To reach the 
phenolphthalein end-point 20 cc. of the acid were required. 
Methyl orange was then added and an additional 30 cc. of 
normal hydrochloric acid was used to cause this indicator to 
change color. Calculate the weights of Na 2 C0 3 and NaHC0 3 
in the mixture. 

106. Write the equations for the production of potassium 
chlorate from potassium chloride by the electrolytic method. 
Calculate the weight of the chlorate which can theoretically be 
obtained from a kilogram of sylvite by this method. What 
volume of hydrogen would be obtained as a by-product? 

107. Calculate the normality of concentrated ammonium 
hydroxide which has a specific gravity of 0.90 and is 28 per 


156 


GENERAL CHEMISTRY 


cent ammonia. Calculate the volume of ammonia gas which 
would be required to prepare a liter of this solution. 

108. Starting with marble give equations and brief statements 
to show how the following could be obtained: (a) quicklime, 
( b ) limewater, (c) calcium bicarbonate, ( d ) calcium chloride, 
( e ) metallic calcium, (/) gypsum, (g) plaster of Paris. 

109. One gram of calcite gave 210 cc. of carbon dioxide 
(standard conditions) when treated with an acid. What was the 
percentage of calcium carbonate in the sample of calcite? 

110. The addition of barium chloride solution to 250 cc. of 
dilute sulfuric acid produced a precipitate which weighed 3.90 g. 
Calculate the normality of the acid. 

111. From solubility relations explain how the following re¬ 
sults are possible: (a) The filtrate from BaC0 3 gives a pre¬ 
cipitate with potassium dichromate solution and with sulfuric 
acid, the filtrate from CaC0 3 does neither. (6) The filtrate 
from CaS0 4 gives a precipitate with ammonium oxalate solu¬ 
tion, but the filtrate from SrS0 4 does not. 

112. Compare the alkali elements and their compounds with 
those of the alkaline earth family. This may be done in tabular 
form. Suggest a precipitation method for separating calcium, 
strontium, and barium from a mixture which contains also 
salts of sodium, potassium and ammonium. 

113. If given 100 g. of silver coin (90 Ag, 10 Cu) and told to 
return the silver as silver chloride, how would you proceed and 
what weight of silver chloride should you return? 

114. An electric current flows through three cells in series: 
Cell A contains a solution of cupric sulfate; B, a solution of 
silver nitrate; and C, dilute sulfuric acid. What weights of 
copper and silver are deposited respectively in A and B while 
3 1. of a mixture of hydrogen and oxygen are being evolved 
from C? 

115. In a Daniell cell what weight of copper has been 
deposited when 5 g. of zinc have dissolved? 

116. Write the equation by which magnesia could be dis¬ 
solved, precipitated and ignited to form magnesium pyrophos¬ 
phate. If a gram sample of insulating material when treated 
in this manner gave 2.35 g. of magnesium pyrophosphate, cal¬ 
culate the percentage of magnesia in the material. 

117. Calculate the volume of carbon dioxide which may be 
obtained from a kilogram of dolomite. Show by equations how 


CHEMICAL ARITHMETIC 


157 


you could start with dolomite, dissolve it, remove the calcium as 
the carbonate, and obtain metallic magnesium from the filtrate. 

118. Write the equations by which zinc may be obtained 
from zinc blende. What weight of zinc and what weight of 
sulfuric acid (by-product) can be obtained from a ton of ore 
which is 30 per cent ZnS? 

119. Calculate the weight of aluminum which should be 
added to 16 g. of ferric oxide to make a thermite mixture. 
Calculate the number of calories liberated when this amount 
of the mixture reacts. 

120. Outline the tests by which you would distinguish the 
following from each other: Aluminum sulfate, zinc sulfate, 
cadmium sulfate, mercuric sulfate, mercurous sulfate. 

121. Construct a series of equations by which bauxite may 
be converted to alum. What weight of alum may be produced 
from a metric ton of bauxite if the ore contains 50 per cent 
A1 2 0 3 ? 

122. By means of ionic equilibria show what happens when a 
solution of sodium carbonate is added to a solution of aluminum 
sulfate. 

123. Starting with tin in each case, write equations to show 
how the following may be prepared: (a) Sodium stannite; 
(i b ) sodium stannate; (c) ammonium sulfostannate. 

124. The sulfides PBS, CuS, HgS, SnS when precipitated are 
dark brown to black in color. Briefly outline chemical tests by 
which each could be distinguished from the others. 

125. Calculate the volume of chlorine, measured at 20° C., 
and 740 mm. pressure, which is required to react completely 
with 2 g. of tin. 

126. Outline the process by which lead is obtained from 
galena. What volume of concentrated sulfuric acid (sp. gr. = 
1.84; 95 per cent H 2 S0 4 ) could be manufactured from the sulfur 
dioxide evolved when a kilogram of galena is roasted? 

127. What chemical test would you make to decide whether 
Pb0 2 is a dioxide or a peroxide? Explain. 

128. Construct the balanced equation for the oxidation of 
ferrous sulfate by potassium dichromate in presence of excess 
sulfuric acid. Calculate the weight of potassium dichromate 
which must be present in a liter of solution so that 1 cc. of the 
solution is equivalent to (will oxidize) 10 mg. of ferrous iron. 

129. Starting with chromic hydroxide show by equation how 


158 


GENERAL CHEMISTRY 


you could obtain: (a) a green solution; (6) a yellow solution; 
(c) a yellow precipitate; ( d ) an orange solution. 

130. In order to separate chromium and aluminum the 
following procedure may be used: (a) To a solution contain¬ 
ing the two metals as chlorides add NaOH; ( b ) add NaOH 
and Na 2 0 2 ; (c) acidify; ( d ) add NH 4 OH. Write a pair of 
equations for each step and indicate how the two metals may 
be separated. 

131. Calculate the volume of sulfur dioxide required to 
reduce 10 g. of potassium dichromate in the presence of excess 
sulfuric acid. Compare the volume of hydrogen sulfide re¬ 
quired for the same purpose. 

132. A quantity of solid potassium dichromate was heated 
with concentrated hydrochloric acid and the gas which was 
produced was passed into a solution of potassium iodide; if 
2.4 g. of iodine were liberated, what weight of potassium dichro¬ 
mate must have been used? 

133. A gram of pyrolusite ore was treated with hydrochloric 
acid and the evolved gas liberated 2.1 g. of iodine from an iodide. 
Calculate the per cent of manganese dioxide in the ore. 

134. Show by equations how the following substances may 
be prepared from manganese dioxide: (a) manganese; ( b ) ferro¬ 
manganese; (c) potassium manganate; ( d ) potassium perman¬ 
ganate; ( e ) manganous sulfate. 

135. What volume of potassium permanganate solution, con¬ 
taining a gram mole of the salt per liter, is required to oxidize 
50 cc. of N/4 sulfurous acid? 

136. An oxide of iron contains 27.6 per cent oxygen; what is 
its formula? If an ore contains 90 per cent of this oxide and 
10 per cent silica, write equations to show how pure iron may 
be obtained from the ore. 

137. How would you prepare a solution of potassium per¬ 
manganate so that 1 cc. of the solution is equivalent to 10 mg. 
of ferrous iron? 

138. How would you prepare: ( a ) ferric chloride from iron; 
(6) ferrous sulfide from ferric chloride; (c) ferric sulfate from 
ferrous sulfide? Write the equations. 

139. Given a mixture of ferric chloride and aluminum chlo¬ 
ride, how could the metals be separated by means of a single 
reagent? Write the equations. Repeat the problem for a 
mixture of ferrous sulfate and copper sulfate. 


CHEMICAL ARITHMETIC 


159 


140. What weight of pure iron will be dissolved by 2 1. of 
6 N sulfuric acid? Calculate the volume occupied by the 
liberated hydrogen at 25° C. and 780 mm. pressure. 

141. If 50 cc. of ferrous sulfate solution, containing excess 
sulfuric acid, reduced 1.58 g. potassium permanganate, what 
was the normality of the solution with respect to ferrous 
sulfate? 

142. Give the equations and the practical details for the 
preparation of pure crystals of ferrous ammonium sulfate from 
iron pyrites. 

143. Starting with ferrous chromite give the equations and 
details by which you would get the iron and chromium into 
separate solutions. 

144. Pyrolusite ore contains ferric oxide and silica as im¬ 
purities. Outline a procedure by which you could obtain a 
pure manganese compound from the ore. 

145. If 10 g. of pure ferric oxide are converted to a solution 
of the sulfate, what volume of concentrated ammonium hy¬ 
droxide (sp. gr. = 0.90; 28 per cent NH 3 ) must be added to 
form ferric ammonium alum? 

146. A solution gave a white precipitate with HC1; the 
precipitate turned black when NH 4 OH was added. The fil¬ 
trate gave a black precipitate with H 2 S which was insoluble in 
HNO 3 but soluble in aqua regia. 

The solution in aqua regia, treated with SnCl 2 , gave a white 
precipitate which changed to gray. The filtrate from the black 
precipitate gave a flocculent white precipitate when treated 
with (NH 4 ) 2 S. This precipitate was soluble in NaOH but not 
appreciably soluble in NH 4 OH. What ions were present in the 
original solution? Write equations for the reactions which 
occurred. 

147. Calculate the weight of impure sodium nitrate, contain¬ 
ing 85 per cent NaN0 3 , required to produce 5 liters of nitric 
acid of specific gravity 1.45 and containing 77 per cent HN0 3 
by weight. 

148. A solution of nitric acid of specific gravity 1.25 contains 
39.82 per cent HN0 3 by weight. Calculate the normality of the 
solution and the weight of HN0 3 in 10 cc. of the solution. 

149. The analysis of a sample of ferric oxide showed that in 
2.5 g. of the oxide there were 1.72 g. of iron. Calculate the 
percentage purity of the ferric oxide. 


160 


GENERAL CHEMISTRY 


150. A solution of potassium hydroxide containing 40 per 
cent by weight of KOH has a specific gravity of 1.4. Calculate 
the weight of potassium chlorate which may be obtained by 
passing an excess of chlorine gas into 600 cc. of this solution. 


APPENDIX 


THE METRIC SYSTEM 


Units 

Meter Liter Gram 

Prefixes Meaning 

Kilo. 1000 

Hecto. 100 

Deka. 10 

Deci. .1 

Centi. .01 

Milli.001 



RELATION OF METRIC AND ENGLISH 
UNITS 



— m 

~ 05 


to — 


1 meter 

= 39.3709 inches 

— 

E 

1 kilogram 

= 2.2046 pounds (Av.) 

- 

M - 

o — 

1 gram 

= 15 grains 

— 


1 liter 

= 1.0567 liquid quarts (U. S.) 

— 

m rz 

H* — 

1 liter 

= 0.9081 dry quart 

— 

~i= 

1 liter 

= 0.8798 Imperial quart (Brit.) 

I_ 

10 — 

1 mile 

= 1.61 kilometers 

- m 

E 

1 pound (Av.) 

= 454 grams 


CO ■ — 

1 pound (Tr.) 

= 373 grams 


»-• — 


— 


- 05 



Fig. 29 


161 


CENTIMETERS 











ELECTROMOTIVE SERIES 


These metals 
NOT 

precipitated 
as sulfides. 


These metals 
precipitated 
as sulfides 
in alkaline 
solution 
(Except Cr) 


Potassium 

Sodium 

Barium 

Strontium 

Calcium 


Magnesium 

Aluminum 

Manganese 

Zinc 

Chromium 
- Cadmium 
Iron 
Cobalt 
Nickel 

\ 

Tin 

Lead 


ui 



Q- 


Hydrogen 


These metals 
precipitated 
as sulfides 
in acid 
solution 


Copper 

Arsenic 

Bismuth 

Antimony 

Mercury 

Silver 

Palladium 

Platinum 

Gold 


5 ^ 


violently 


slowly 


162 


THESE METALS LIBERATE HYDROGEN FROM: 









APPROXIMATE ATOMIC WEIGHTS 



Symbol 

Atomic 

Weight 


Symbol 

Atomic 

Weight 

Aluminum . .. 

A1 

27.0 

Iron. 

Fe 

56.0 

Antimony. . . . 

Sb 

121.0 

Lead. 

Pb 

207.0 

Arsenic. 

As 

75.0 

Magnesium. .. 

Mg 

24.0 

Barium. 

Ba 

137.0 

Manganese. . . 

Mn 

55.0 

Bismuth. 

Bi 

208.0 

Mercury. 

Hg 

200.0 

Boron 

B 

11.0 

Nickel. 

Ni 

58.5 

Bromine. 

Br 

80.0 

Nitrogen. 

N 

14.0 

CJalmiim 

Ca 

40.0 

Oxygen. 

O 

16.0 

Carbon. 

C 

12.0 

Phosphorus... 

P 

31.0 

Chlorine. 

Cl 

35.5 

Platinum. 

Pt 

195.0 

Chromium .. . 

Cr 

52.0 

Potassium.... 

K 

39.0 

Cobalt 

Co 

59.0 

Radium. 

Ra 

226.0 

Copper. 

Cu 

63.5 

Silicon. 

Si 

28.0 

TPIn nrinp 

F 

19.0 

Silver. 

Ag 

108.0 

JL iUUIlilV/ . 

Gold 

Au 

197.0 

Sodium. 

Na 

23.0 

TTplinm 

He 

4.0 

Sulfur. 

S 

32.0 

U.111 • ••••• 

Hydrogen.... 

H 

1.0 

Tin. 

Sn 

119.0 

Iodine. 

I 

127.0 

Zinc. 

Zn 

65.0 


VAPOR PRESSURE OF WATER 


Temperatures 

Pressure in mm. 
of Mercury 

Temperature 

Pressure 

IN MM. 

10° 

9.2 

25° 

23.5 

11° 

9.8 

26° 

25.0 

12° 

10.5 

27° 

26.5 

13° 

11.2 

28° 

28.1 

14° 

11.9 

29° 

29.8 

15° 

12.7 

30° 

31.6 

16° 

13.6 

31° 

33.4 

17° 

14.5 

32° 

35.4 

18° 

15.4 

33° 

37.4 

19° 

16.4 

o 

CO 

39.6 

20° 

17.4 

o 

CO 

41.9 

21° 

18.5 



22° 

19.7 

99° 

733.2 

23° 

20.9 

100° 

760.0 

24° 

22.2 

101° 

787.6 


163 

















































FIRST AID INSTRUCTIONS 


Minor Cuts and Bruises. — Flush wound with water, dry with 
sterile gauze, apply mercurochrome solution (use sparingly), 
bandage with roller gauze. 

Severe Bleeding. — If from a cut artery, stop with pledget of 
gauze directly on leak or apply a tourniquet (towel or rubber 
tube) between wound and heart. Coagulation of blood may 
be hastened by application of pledget of gauze moistened 
with aluminum acetate. Bandage with gauze and roller 
bandage. CAUTION, NEVER APPLY COTTON TO AN 
OPEN WOUND. 

Bums, Steam and Hot Water. Apply cold running water at 
once. When affected parts have been cooled remove cloth¬ 
ing at your leisure, dry, apply butesin picrate and bandage 
loosely if necessary. 

Burns, Fire, Hot Objects, etc. — Apply butesin picrate and 
bandage if necessary. 

Bums by Chemicals. 

Acids. — Flood with water, apply saturated solution of 
sodium bicarbonate. 

Alkalies. — Flood with water. Apply boric acid solution 
or 1 to 2 per cent acetic acid. 

Bromine. — Flood with water. Apply sodium bicarbonate 
solution or sodium thiosulfate solution. Keep wound 
moist with glycerine. 

Phosphorus. — Flood with water. Apply mixture of 
sodium bicarbonate and hydrogen peroxide. 

Poisons Inhaled as Gases and Fumes — General Instructions. 
— Secure fresh air and have patient breathe deeply. Do not 
stimulate circulation by rubbing and chafing. 

Acid Vapors. — Counteract with inhalation of dilute am¬ 
monia. 

Hydrogen Sulfide Gas. — Give fresh air. Inhale ammonia. 
Chlorine and Bromine Gas. — Give fresh air. Inhale vapors 
of alcohol or, better still, inhale vapors of tincture of 
benzoin. 

Fainting. — Secure fresh air, recline patient, inhale ammonia 
(smelling salts). In more severe cases heart action may be 
stimulated by aromatic spirits of ammonia. 

164 


y 

# 

























